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The role of N2O and NNH in the formation of NO via HCN in hydrocarbon flames
Combustion and Flame 133 (2003) 311–322
The role of N2O and NNH in the formation of NO via HCN in hydrocarbon flames J. Tomeczek*, B. Gradon´ Department of Process Energy, Silesian Technical University, ul.Krasin´skiego 8, 40-019 Katowice, Poland Received 14 April 2002; received in revised form 2 January 2003; accepted 14 January 2003
Abstract A new way of forming HCN in flames via N2O and NNH reacting with CHi radicals is proposed and tested for rich and lean gaseous premixed flames of CH4 and air and also of CH4, N2O and Ar. This new route is thermodynamically more probable than Fenimore’s direct reaction of N2 with CHi radicals. In fact, it is shown that the new mechanism is more important than Fenimore’s reaction in both rich and lean flames. Rate constants of the new reactions forming NO have been suggested on the basis of numerical modeling. It has been shown that the formation of NO through HCN is most effective as the result of reactions initiated by N2O ⫹ CH3 3 CH2NH ⫹ NO, followed by CH2NH ⫹ H 3 H2CN ⫹ H2 and CH2NH ⫹ O 3 H2CN ⫹ OH. In flames of CH4 and air, a substantial source of N2O comes from the reverse of the reaction N2O ⫹ CH3 7 CH3O ⫹ N2 in the reaction zone. A formula based on the steady state assumption and partial equilibrium limits the number of nitrogen conversion reactions to only 12; this was tested using a premixed flame of CH4 and air. © 2003 The Combustion Institute. All rights reserved. Keywords: HCN; N2O; NNH and NO formation
Introduction Modeling [NO] in gaseous hydrocarbon flames still creates fundamental difficulties. There are two distinct problems: the details of the mechanism of chemical reactions producing NO, together with values of their rate constants; the concentration of radicals participating in this mechanism, mainly [O], [H] and [CHi]. Zeldovich [1] proposed a formation mechanism of two reactions producing thermal NO in: N2 ⫹ O 3 NO ⫹ N,
(1)
N ⫹ O2 3 NO ⫹ O,
(2)
and calculated [O] assuming that the thermal disso* Corresponding Author. Tel./Fax: ⫹48-32-603-4286. E-mail address: [email protected] (J. Tomeczek).
ciation of O2 in O2⫹ M N O ⫹ O ⫹ M was at equilibrium. This assumption enabled him to find the rate constants of the thermal NO mechanism. This approach was fully justified, because the flame temperature in Zeldovich’s experiments was above 2000°C. In many flames, however, [O] is much higher than [O]eq calculated from thermal equilibrium, but even using the real [O], it is not possible to explain the observed [NO] in hydrocarbon and other flames using Zeldovich’s reactions (1) and (2). Fenimore [2] accordingly proposed the so-called prompt reaction between molecular nitrogen N2 and CHi radicals in hydrocarbon flames. For the Fenimore’s mechanism, overall rate equations were proposed [3,4], which often were based on subtracting from the observed rate of NO the rate of Zeldovich’s reactions (1) and (2) calculated using [O]eq. Such equations include any reaction taking place within the flame and not only Fenimore’s reaction, which
0010-2180/03/$ – see front matter © 2003 The Combustion Institute. All rights reserved. doi:10.1016/S0010-2180(03)00013-0
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J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
was extensively studied [5]. Indeed, its contribution was postulated to be important in cooler, rich hydrocarbon flames, whereas in hotter, lean flames, the amount of prompt NO fell abruptly. Fenimore and Fraenkel [6] later pointed out that the role of the N2 ⫹ CHi had been probably overvalued. Luque et al. [7] found that the rate of Fenimore’s principal reaction N2 ⫹ CH 3 HCN ⫹ N
(56)
should be about twice as fast than traditionally assumed to reproduce the measured [NO] in flames of methane ⫹ air. Wolfrum [8] and Malte and Pratt [9] proposed the nitrous oxide mechanism based on three reactions: N2O ⫹ M7N2 ⫹ O ⫹ M,
(3)
N2O ⫹ O 3 NO ⫹ NO,
(4)
N 2 O ⫹ O 3 N2 ⫹ O2 .
(5)
However, the recommended rate constants of these reactions [10] limited the contribution of this mechanism to only ⬃12% of the rate of formation of NO. Tomeczek and Gradon´ [11] found that the thermal mechanism is much faster than described by the Zeldovich reactions with the recommended rate constants. They explained it by introducing the term “extended thermal mechanism” including both Zeldovich’s reactions (1 and 2) and the nitrous oxide reactions (3 to 5), for which the rate constants of the controlling reactions were proposed and substantiated with experimental evidence from the literature. The extended thermal mechanism can reproduce the observed [NO] in lean CH4 ⫹ air flames and in rich H2 ⫹ CO ⫹ N2O ⫹ Ar flames [12] if proper [O] are used. However, for rich flames of CH4 ⫹ air, the calculated [NO] are still too small; this suggests that further reactions operate in these conditions. Bozzelli and Dean [13] introduced the NNH mechanism composed of the two main reactions: NNH ⫹ M7N2 ⫹ H ⫹ M, NNH ⫹ O 3 NH ⫹ NO,
(9) (10)
for which they proposed the kinetic constants. They came also to the conclusion that within a flame the first reaction is close to equilibrium. The aim of this research was to tackle the modeling of [NO] in rich and near-stoichiometric flames of CH4 ⫹ air using available experimental data for gaseous fuels. However, rich flames do not occur often. Nevertheless, in practical diffusion flames, there are large volumes containing rich mixtures where most of the NO can be formed. The difficulty has been that any additional mechanism must be able
to increase [NO] in rich conditions, but at the same time to give less NO when the flame’s composition shifts toward lean conditions. It has been found that these difficulties can be solved using the new proposed reactions of CHi radicals with N2O and NNH. These, together with the extended thermal mechanism and the modified NNH mechanism, predict very well [NO] in hydrocarbon flames. A simplified version of the mechanism is proposed also for modeling large practical flames and is tested with a flame of CH4 ⫹ air.
Proposed mechanism for the formation of HCN Following Fenimore [2], a direct attack of a hydrocarbon radical on molecular N2 is traditionally considered as the first step on the road to NO via HCN as an intermediate [5]. However, the reactions of CHi with N2 are endothermic with positive values of the change in Gibbs free energy ⌬Go, thus giving very small equilibrium constants for these reactions, so the contribution to [NO] must be insignificant in most flame conditions. Blauwens et al. [14] on the basis of mass spectrometric measurements found that it is impossible to distinguish between reactions (56) and (57). Miller and Bowman [10] concluded that only Fenimore’s principal reaction (56) cannot be ruled out; however, they also included reaction (57) in their system of reactions for producing NO. The shortcomings of Fenimore’s mechanism are possibly overcome by the NNH mechanism; however, the results must be taken cautiously. Hayhurst and Hutchinson [15] could not find any evidence for the production of NO via N2O, but confirmed that significant amounts of NO are formed via NNH through reaction (10) in rich flames of H2 ⫹ O2 ⫹ N2. This has been tested by Konnov et al. [16]. As a matter of fact, the difference between the experimental [NO] and those modeled by Miller and Bowman’s [10] mechanism were taken [15] as a contribution to reaction (10), with subsequent total conversion of NH into NO. It is necessary to emphasize that most of these difficulties can be explained using the extended thermal mechanism proposed by Tomeczek and Gradon´ [11], where N2O plays a very important role. What is more, there is direct experimental evidence, gathered from the literature by Tomeczek and Gradon´ [11], that the rate of reaction (4) is much faster than traditionally assumed, while for the reaction (10) the rate constant was not measured directly. Even in a flow of air without fuel through a high temperature reactor, the measured [NO] at the outflow will be from 3 to 5 times higher [11] than predicted by the Zeldovich mechanism with Miller and Bowman’s [10] rate constants. The NNH mech-
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
anism with the available rate constants together with the extended thermal mechanism [11] considerably overpredicts [NO] in lean CH4 ⫹ air flames [12], which suggests that the rate coefficient of reaction (10) proposed by Bozzelli and Dean [13] and Hayhurst and Hutchinson [15], recently corrected by Konnov and de Ruyck [17], is too high. The presence of HCN in flames with a hydrocarbon as fuel is an experimental fact, but in addition to (56) and (57) it can also be formed [18] exothermically by hydrocarbon (CHi) fragments reacting with N2O and NNH. The elementary steps of these reactions are discussed below.
flames should be expected for (66). Nevertheless, it has been included in the final mechanism, but with a low rate coefficient. The interesting feature of reaction (67) is that in some flame regions it can produce N2O as a result of high [N2] and [CH3O]. This reaction is exothermic; nevertheless, in regions of high [CH3O], thermodynamics (⌬G0 ⫽ ⫺218.3 kJ/mol) favors its backwards step. The newly proposed reverse of reaction (67) increases [N2O] so much that N2O can become a dominant source for HCN. The subsequent reactions most probably leading to the formation of HCN are: CH2NH ⫹ H 3 H2CN ⫹ H2, ⌬H ⫽ ⫺ 54.6 kJ/mol, 0
Discussion
313
(68) (68)
CH2NH ⫹ O 3 H2CN ⫹ OH, The state of the art leads one to expect that the role of NNH in the formation of HCN should be much more important than that of N2O, because of the high rate of formation [13] of NNH. The introduction of new reactions consuming NNH should not influence its concentration due to [NNH] being close to equilibrium in flames [13]. On the other hand, the reactions of CH3 radicals are most likely to produce enough HCN in flames, because [CH] and [CH2] are lower than [CH3]. The reaction of NNH with CH3 can proceed to a known products in: NNH ⫹ CH3 3 N2 ⫹ CH4, ⌬H ⫽ ⫺471.3 kJ/mol, 0
(64)
N2O ⫹ CH3 3 CH2NH ⫹ NO, (65)
NNH ⫹ CH3 3 CH2N2 ⫹ H2, ⌬H0 ⫽ ⫺ 110.3 kJ/mol.
(66)
However the reactants of (65) can also undergo the parallel reaction: N2O ⫹ CH37CH3O ⫹ N2, ⌬H0 ⫽ ⫺ 212.3 kJ/mol.
(69)
CH2N2 ⫹ O 3 H2CN ⫹ NO, ⌬H0 ⫽ ⫺ 196.9 kJ/mol,
(71)
H2CN ⫹ H 3 HCN ⫹ H2, ⌬H0 ⫽ ⫺ 334.7 kJ/mol,
(72)
H2CN ⫹ O 3 HCN ⫹ OH, ⌬H0 ⫽ ⫺ 326.5 kJ/mol,
(73)
CH2N2 ⫹ O 3 HCNN ⫹ OH,
which does not yield NO. It is also possible that the same reactants initially form the stable molecule N2CH4 [19] in: NNH ⫹ CH3 ⫹ M 7 N2CH4 ⫹ M, which must have an unrealistically high rate coefficient in order to contribute to the formation of HCN. An alternative way to HCN is through the known molecules CH2N2 and CH2NH [20], which can be formed in:
⌬H0 ⫽ ⫺53.3 kJ/mol.
⌬H0 ⫽ ⫺ 46.4 kJ/mol,
(67)
The four-centered reaction (66) is similar in character to Fenimore’s second reaction (57); thus, a minor role in the formation of NO in premixed hydrocarbon
⌬H0 ⫽ ⫺ 33.9 kJ/mol.
(74)
Thus, the H2CN formed in reactions (68) and (69) may be converted into HCN in reactions (72), (73), and (76), while the HCNN formed in reaction (74) may give HCN in reaction (100). The rate constants of these reactions can only be suggested, because at present no reliable experimental meassurements are available. Thus, a numerical, modeling way has been chosen to adjust the rate constants. The published chemical compositions of the flames of CH4 ⫹ air [15,21,22] do allow for the rate coefficients to be determined, but only with an activation energy of zero, which is controversial. This problem of the activation energy could be partially solved using the detailed measurements of the flame of CH4 ⫹ N2O ⫹ Ar by Vandooren et al. [23], because of a broad range and variation of stable species and radicals in this flame. However, a scheme of only the above reactions was unable to describe the composition of this flame, which was very difficult to model. Particular problems were to compute the concentrations of HNCO, HCN, and NO. The known reactions involving HNCO (107-122) could not solve the problem; thus the following new reactions facilitated the modeling:
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N2O ⫹ HCN 3 HNCO ⫹ N2, ⌬H0 ⫽ ⫺330.7 kJ/mol,
(104)
HNCO ⫹ NO 3 NNH ⫹ CO2, ⌬H0 ⫽ ⫺117.2 kJ/mol,
(105)
HNCO ⫹ NH27NCO ⫹ NH3, ⌬H0 ⫽ ⫹12.0 kJ/mol.
(106)
When formulating rate coefficients, the pre-exponential factor was kept a safe margin below the upper limit imposed by kinetic theory. For all the proposed reactions, the suggested pre-exponential factors and activation energies are given in Table I. The crucial reaction forming HCN in the proposed mechanism is (65). However, without the reverse of reaction (67) as the source of N2O, (65) is not effective. Both these reactions must be fast, otherwise the measured profiles of [N2O] and [CH3] along Vandooren et al.’s [23] flame cannot be described. As far as reaction (65) is concerned, there is no previous work for comparison; however, for reaction (67) there are two reports suggesting that at 873 K [24,25], this reaction is slow with a rate coefficient of ⬃14 m3/(mol s). These results were obtained from simplified modeling involving only 11 reactions; in fact, many of the species measured in large quantities by Vandooren et al. [23] were not included. Using the rate coefficient of reaction (67) assumed by Borisov et al. [25], the experiment by Vandooren et al. [23] could not be described. Our finding was that the rate coefficients suggested in Table I provided the only solution to match experimental evidence. Otherwise, Miller and Bowman’s [10] combustion mechanism was assumed. Numerical simulations for rich CH4 ⫹ air flames have shown that HCN is formed in the new reactions earlier than in Fenimore’s principal reaction (56), because CH3 precedes CH in such a flame. The contribution of the proposed reactions was tested for a rich mixture (air ⫽ 0.9 stoichiometric amount) of CH4 ⫹ air of constant temperature 1800 K. The results are presented in Fig. 1. The new mechanism generates more HCN than Fenimore’s reactions, for which the [HCN] profile is very narrow. However, the maximum in [HCN] for both mechanisms is observed in similar positions, so [NO] predicted by the new mechanism is also higher than described by Fenimore’s mechanism. The CH4 flames analyzed by Malte et al. [21], Sarofim and Pohl [22], and by Hayhurst and Hutchinson [15] were chosen to verify the modeling of hydrocarbon flames. Altogether, four different mechanisms were analyzed: three (Miller and Bowman [10], extended thermal mechanism (1 to 5); [11]
extended thermal mechanism (1 to 5) plus reaction (6) from Bozzelli et al. [23]) reflected current knowledge, but the fourth new one comprised reactions (1 to 122) from Table I. Figure 2 presents the results for three cases of modeling using available mechanisms in the literature for the experimental measurements of Malte et al. [21]. In the modeling, the flame temperature was assumed to be the mean of the values measured by a thermocouple and by the OH rotational method. Miller and Bowman’s [10] mechanism produces [NO] much lower than the experimental values for both the lean and rich flames. The extended thermal mechanism is also not satisfactory, even after including reaction (6). For very lean flames, the extended thermal mechanism is able to account for over 50% of the observed [NO]. The [NO] generated in a slightly rich flame of CH4 ⫹ air as investigated by Sarofim and Pohl [22] are presented in Fig. 3, together with the modeling results for the three mechanisms representing the state of art. Again, the extended thermal mechanism with the additional reaction (6) produces reasonable agreement with the experimental points. For the very rich CH4 ⫹ air flame of Hayhurst and Hutchinson [15], the extended thermal mechanism with reaction (6) also gives the smallest divergence from the experimental points as seen in Fig. 4. However, still over 20% of the [NO] is not predicted. None of the three mechanisms analyzed using known rate constants can describe satisfactorily the observed [NO]. This suggests that some important reactions have been neglected. To be objective, the consequences of modifying of the Fenimore’s reactions by enhancing of their rate constants were examined. In this way it was possible to improve the modeling in rich flames, but unfortunately in flames near to stoichiometric and in lean ones the modeled of [NO] were too high. Assessment of reaction (10) has to take into account the additional contribution of the new reactions of N2O and NNH with CHi radicals. A proper value of the (10) reaction rate constant can be found only using hydrogen flames to exclude the contribution of HCN. The difference between the experimental [NO] values and those calculated by the extended thermal mechanism with the rate constants by Tomeczek and Gradon´ [11] has been taken as the contribution of reaction (10) in hydrogen flames; the evaluated rate constants for reaction (10) are given in Table I, and are smaller than those proposed by Bozzelli and Dean [13], because they used Miller and Bowman’s [10] rate constants for the thermal mechanism. The final consequence for modeling NO is demonstrated successfully in Fig. 5 for a very rich H2 ⫹ air flame investigated by Hayhurst and Hutchinson [15]. Three
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
315
Table I Rate constants of nitrogen conversion reactions, kf ⫽ k0 Tn exp(⫺E/RT) No. 1 2 3 Enhanced 4 5 6 7 8 9 Enhanced 10 11 12 13 14 15 16 17 18 19 20 Enhanced 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 Enhanced 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
2.72 ⫻ 108 N2⫹O7NO⫹N 0.00 6.40 ⫻ 103 N⫹O27NO⫹O 1.00 3.20 ⫻ 108 N2O⫹M7N2⫹O⫹M 1.00 third-body efficiencies: O2 (1.4), CO2 (3.0), N2 (1.7), H2O (12.0), other (1.0) N2O⫹O7NO⫹NO 6.21 ⫻ 108 0.00 N2O⫹O7N2⫹O2 6.21 ⫻ 108 0.00 2.50 ⫻ 1010 NH⫹NO7N2O⫹H ⫺1.03 2.20 ⫻ 108 N2O⫹H7N2⫹OH 0.00 N2O⫹OH7N2⫹HO2 1.64 ⫻ 105 0.00 1.30 ⫻ 108 ⫺0.11 NNH⫹M7N2⫹H⫹M third-body efficiencies: all species (1.0) NNH⫹O7NO⫹NH 3.30 ⫻ 107 ⫺0.23 2.90 ⫻ 105 NNH⫹O27N2O⫹OH ⫺0.34 1.40 ⫻ 108 NNH⫹O7N2O⫹H ⫺0.40 NNH⫹H7N2⫹H2 1.00 ⫻ 105 0.00 2.40 ⫻ 1016 NNH⫹OH7N2⫹H2O ⫺2.88 1.70 ⫻ 1010 NNH⫹O7N2⫹OH ⫺1.23 NNH⫹O27N2⫹HO2 1.20 ⫻ 106 ⫺0.34 5.00 ⫻ 107 NNH⫹NO7N2⫹HNO 0.00 NNH⫹NH7N2⫹NH2 5.00 ⫻ 107 0.00 NNH⫹NH27N2⫹NH3 5.00 ⫻ 107 0.00 1.10 ⫻ 1010 NO2⫹M7NO⫹O⫹M 0.00 third-body efficiencies: all species (1.0) 8.43 ⫻ 107 NO2⫹H7NO⫹OH 0.00 NO2⫹O7NO⫹O2 3.91 ⫻ 106 0.00 2.11 ⫻ 106 NO⫹HO27NO2⫹OH 0.00 7.60 ⫻ 104 NH⫹O27NO⫹OH 0.00 NH⫹O7NO⫹H 5.50 ⫻ 107 0.00 5.00 ⫻ 105 NH⫹OH7N⫹H2O 0.50 3.00 ⫻ 107 NH⫹N7N2⫹H 0.00 NH⫹O7N⫹OH 3.72 ⫻ 107 0.00 2.16 ⫻ 107 NH⫹NO7N2⫹OH ⫺0.23 1.60 ⫻ 108 N⫹H27NH⫹H 0.00 5.10 ⫻ 107 NH⫹NH7N2⫹H⫹H 0.00 NH⫹OH7NO⫹H2 2.00 ⫻ 107 0.00 3.89 ⫻ 107 NH⫹O27HNO⫹O 0.00 NH⫹OH7HNO7H 2.00 ⫻ 107 0.00 HNO⫹M7NO⫹H⫹M 1.50 ⫻ 1010 0.00 third-body efficiencies: H2O (10.0), O2 (2.0), N2 (2.0), H2 (2.0), other (1.0) HNO⫹OH7NO⫹H2O 4.80 ⫻ 107 0.00 HNO⫹O7NO⫹OH 3.61 ⫻ 107 0.00 1.81 ⫻ 107 HNO⫹H7H2⫹NO 0.00 3.95 ⫻ 106 HNO⫹HNO7N2O⫹H2O 0.00 2.00 ⫻ 106 HNO⫹NO7N2O⫹OH 0.00 N⫹OH7NO⫹H 3.80 ⫻ 107 0.00 7.00 ⫻ 106 NH2⫹O7OH⫹NH 0.00 6.63 ⫻ 108 NH2⫹O7HNO⫹H ⫺0.50 4.00 2.00 NH2⫹OH7NH⫹H2O NH2⫹NO7NNH⫹OH 6.40 ⫻ 109 ⫺1.25 6.20 ⫻ 1012 NH2⫹NO7N2⫹H2O ⫺1.25 7.20 ⫻ 107 NH2⫹N7N2⫹H⫹H 0.00 4.50 ⫻ 106 NH2⫹O27HNO⫹OH 0.00 NH2⫹H7NH⫹H2 6.92 ⫻ 107 0.00 2.00 ⫻ 107 NH2⫹HNO7NH3⫹NO 0.00 1.40 ⫻ 1010 NH3⫹M7NH2⫹H⫹M 0.00
E kJ mol⫺1
References
318.59 26.19 232.13
[11] [10] [30]
108.81 108.81 3.49 70.10 41.57 20.84
[11] [11] [23] [23] [30] [13]
⫺4.23 0.62 1.99 0.00 10.23 2.08 0.62 0.00 0.00 0.00 274.36
[18] [13] [13] [13] [13] [13] [13] [10] [10] [10] [10]
0.00 1.00 ⫺2.01 6.41 0.00 8.37 0.00 0.00 0.00 105.26 0.00 0.00 74.88 0.00 203.82
[31] [31] [10] [10] [38] [10] [10] [38] [39] [32] [32] [40] [38] [10] [10]
4.14 0.00 4.16 20.93 108.86 0.00 0.00 0.00 4.19 0.00 0.00 0.00 104.67 15.28 4.19 379.34
[31] [31] [31] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [27] continued
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Table I Continued No.
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
Enhanced third-body efficiencies: all species (1.0) 5.00 ⫻ 1010 NH3⫹M7NH⫹H2⫹M 0.00 Enhanced third-body efficiencies: all species (1.0) 53 NH3⫹H7NH2⫹H2 0.64 2.39 2.04 2.04 54 NH3⫹OH7NH2⫹H2O 55 NH3⫹O7NH2⫹OH 2.10 ⫻ 107 0.00 3.00 ⫻ 105 56 N2⫹CH7HCN⫹N 0.00 1.00 ⫻ 107 57 N2⫹CH27HCN⫹NH 0.00 5.00 ⫻ 107 58 N⫹CH37HCN⫹H⫹H 0.00 0.00 59 NO⫹CH7HCN⫹O 1.10 ⫻ 108 1.00 ⫻ 105 60 NO⫹CH37HCN⫹H2O 0.00 2.00 ⫻ 107 61 NO⫹CH27HCN⫹OH 0.00 8.00 ⫻ 107 62 N2O⫹CH7HCN⫹NO 0.00 63 NNH⫹CH7HCN⫹NH 8.00 ⫻ 107 0.00 64 NNH⫹CH37N2⫹CH4 2.50 ⫻ 107 0.00 6.73 ⫻ 104 65 N2O⫹CH37CH2NH⫹NO 2.00 66 NNH⫹CH37CH2N2⫹H2 1.00 ⫻ 107 0.00 67 N2O⫹CH37CH3O⫹N2 1.90 ⫻ 104 1.75 68 CH2NH⫹H7H2CN⫹H2 8.32 ⫻ 103 1.25 1.62 ⫻ 104 69 CH2NH⫹O7H2CN⫹OH 1.25 4.00 ⫻ 107 70 CH2NH⫹O7CH2O⫹NH 0.00 8.00 ⫻ 107 71 CH2N2⫹O7H2CN⫹NO 0.00 72 H2CN⫹H7HCN⫹H2 5.00 ⫻ 107 0.00 5.00 ⫻ 107 73 H2CN⫹O7HCN⫹OH 0.00 2.00 ⫻ 107 74 CH2N2⫹O7HCNN⫹OH 0.00 75 H2CN⫹N7N2⫹CH2 6.00 ⫻ 107 0.00 3.30 ⫻ 101 76 HCN⫹H⫹M7H2CN⫹M 0.00 Enhanced third-body efficiencies: H2 (2.0), H2O (6.0), CH4 (2.0), CO (1.5), CO2 (2.0), C2H6 77 CH3⫹NO7H2CN⫹OH 1.00 ⫻ 106 0.00 6.10 ⫻ 108 78 CH3⫹N7H2CN⫹H ⫺0.31 79 HCN⫹O7NH⫹CO 3.45 ⫻ 10⫺3 2.64 80 HCN⫹O7NCO⫹H 1.38 ⫻ 10⫺2 2.64 7.83 ⫻ 10⫺10 81 HCN⫹OH7NH2⫹CO 4.00 82 NCO⫹H7NH⫹CO 5.00 ⫻ 107 0.00 83 NCO⫹O7NO⫹CO 2.00 ⫻ 107 0.00 2.00 ⫻ 107 84 NCO⫹N7N2⫹CO 0.00 85 NCO⫹M7N⫹CO⫹M 3.10 ⫻ 1010 ⫺0.50 Enhanced third-body efficiencies: N2 (1.5), O2 (1.5), H2O (18.6), other (1.0) 86 NCO⫹NO7N2O⫹CO 1.00 ⫻ 107 0.00 87 NCO⫹OH7NO⫹CO⫹H 1.00 ⫻ 107 0.00 1.90 ⫻ 105 88 N⫹CO27NO⫹CO 0.00 89 CH⫹N7CN⫹H 1.30 ⫻ 107 0.00 1.45 ⫻ 107 90 HCN⫹OH7CN⫹H2O 0.00 91 HCN⫹O7CN⫹OH 2.70 ⫻ 103 1.58 2.92 ⫻ 10⫺1 92 CN⫹H27HCN⫹H 2.45 93 CN⫹O7CO⫹N 1.80 ⫻ 107 0.00 5.60 ⫻ 106 94 CN⫹O27NCO⫹O 0.00 95 CN⫹OH7NCO⫹H 6.00 ⫻ 107 0.00 3.00 ⫻ 107 96 CN⫹NO27NCO⫹NO 0.00 97 CN⫹N2O7NCO⫹N2 1.00 ⫻ 107 0.00 3.10 0.15 98 CH⫹N2⫹M7HCNN⫹M Enhanced third-body efficiencies: H2 (2.0), H2O (6.0), CH4 (2.0), CO (1.5), CO2 (2.0), C2H6 99 HCNN⫹O7CO⫹H⫹N2 2.20 ⫻ 107 0.00 52
E kJ mol⫺1
References
340.00
[41]
42.59 2.37 37.68 56.94 309.83 0.00 0.00 62.80 0.00 0.00 0.00 0.00 127.00 0.00 80.00 10.00 20.00 20.00 0.00 0.00 0.00 0.00 1.67 0.00 (3.0), other (1.0) 91.07 1.21 20.85 20.85 16.75 0.00 0.00 0.00 201.00
[10] [10] [10] [10] [10] [27] [10] [10] [10] this work this work [43] this work this work this work this work this work this work this work this work this work this work [43] [43]
⫺1.63 0.00 14.24 0.00 45.76 111.37 9.37 0.00 0.00 0.00 0.00 0.00 0.00 (3.0), other (1.0) 0.00
[10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [43]
[43] [43] [10] [10] [10] [10] [10] [10] [42]
[43] continued
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317
Table I Continued No.
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
E kJ mol⫺1
References
100 HCNN⫹O7HCN⫹NO 2.00 ⫻ 106 0.00 101 HCNN⫹O27O⫹CHO⫹N2 1.20 ⫻ 107 0.00 102 HCNN⫹OH7H⫹CHO⫹N2 1.20 ⫻ 107 0.00 103 HCNN⫹H7CH2⫹N2 1.00 ⫻ 108 0.00 104 N2O⫹HCN7HNCO⫹N2 1.20 ⫻ 107 0.00 105 HNCO⫹NO7NNH⫹CO2 2.00 ⫻ 107 0.00 106 HNCO⫹NH27NCO⫹NH3 3.70 ⫻ 106 0.00 2.00 ⫻ 107 107 HNCO⫹H7NH2⫹CO 0.00 108 OH⫹HCN7HNCO⫹H 1.98 ⫻ 10⫺9 4.00 8.59 ⫻ 106 0.00 109 NCO⫹H27HNCO⫹H 110 HNCO⫹M7NH⫹CO⫹M 1.10 ⫻ 1010 0.00 Enhanced third-body efficiencies: N2 (1.5), O2 (1.5), H2O (18.6), other (1.0) 111 HNCO⫹O7NCO⫹OH 2.20 2.11 112 HNCO⫹O7NH⫹CO2 9.60 ⫻ 101 1.41 6.40 ⫻ 10⫺1 113 HNCO⫹OH7NCO⫹H2O 2.00 3.10 ⫻ 1011 114 CH2⫹NO7H⫹HNCO ⫺1.38 115 HNCO⫹O7HNO⫹CO 1.50 ⫻ 102 1.57 116 HNCO⫹OH7NH2⫹CO2 3.30 1.50 3.80 ⫻ 107 117 CH2⫹NO7HCNO⫹H ⫺0.36 118 HCNO⫹H7HCN⫹OH 2.70 ⫻ 105 0.18 1.70 ⫻ 108 119 HCNO⫹H7NH2⫹CO ⫺0.75 120 HOCN⫹H7HNCO⫹H 2.00 ⫻ 101 2.00 121 HCNO⫹H7HNCO⫹H 2.10 ⫻ 109 ⫺0.96 122 HCN⫹OH7HOCN⫹H 2.00 ⫻ 101 2.00
0.00 0.00 0.00 0.00 0.00 0.00 10.00 12.56 4.19 37.68 360.00
[40] [40] [40] [40] thiswork thiswork thiswork [10] [10] [10] [42]
47.90 35.70 10.70 5.32 184.23 15.07 2.43 8.88 12.10 8.37 11.93 8.37
[42] [42] [42] [43] [43] [43] [43] [43] [43] [43] [43] [43]
of the analyzed mechanisms are unable to reproduce the measured [NO], because the calculated [NO] are approximately three times lower than the experimental ones, and only the fourth solid line for the modeling using the extended thermal mechanism plus reaction (10) describes well the experiments’ results. The results of the modeling using proposed mechanism are shown as solid lines in Figs. 2– 4. Thus, in all the flames analyzed, this mechanism
Fig. 1. Comparison of [NO] and [HCN] generated by the Fenimore mechanism (- - -) and as modeled by proposed mechanism via HCN (——).
has proved its ability to reproduce the experimental results for CH4 flames. The contribution of the key proposed reactions into forming NO via the inter-
Fig. 2. Comparison of the [NO] measured by Malte et al. [21] (Œ) within the premixed flames of CH4 ⫹ air and calculated by four mechanisms: (- - -) Miller and Bowman [10], (— 䡠 —) extended thermal mechanism (1-5) [11], (— —) extended thermal mechanism (1-5) plus reaction (26) from Bozzelli et al. [26] and (——) mechanism proposed here.
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Fig. 3. Comparison of [NO] measured by Sarofim and Pohl [22] (Œ) along a slightly rich premixed flame of CH4 ⫹ air (air ⫽ 0.95 ⫻ stechiometric) and modeled by four mechanisms: (- - -) Miller and Bowman [10], (— 䡠 —) extended thermal mechanism (1-5) [11], (— —) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (——)mechanism proposed here.
mediate HCN in rich flames of CH4 ⫹ air is higher than the contribution of Fenimore’s reactions. The extended thermal mechanism produces ⬃15% of the [NO] in this case. In lean CH4 ⫹ air flames, however, the most important contribution is from the extended thermal mechanism, in which the reaction (4) is decisive. In rich CH4 ⫹ air flames the proposed mechanism is the most important source of NO through HCN; in fact, the reactions initiated by N2O with
Fig. 4. Comparison of [NO] measured by Hayhurst and Hutchinsom [15] (Œ) along a rich premixed flame of CH4 ⫹ air (air ⫽ 0.67 ⫻ stoichiometric) and modeled by four mechanisms: (a) proposed mechanism, (b) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (c) Miller and Bowman [10], (d) extended thermal mechanism (1-5) [11].
Fig. 5. Comparison of [NO] measured by Hayhurst and Hutchinsom [15] (Œ) along a very rich premixed flame of H2 ⫹ air (air ⫽ 0.67 ⫻ stechiometric) and modeled by four mechanisms: (a) proposed mechanism, (b) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (c) Miller and Bowman [10], (d) extended thermal mechanism (1-5) [11].
CH3 in (65) are ⬃25 times more effective than the reactions initiated by NNH with CH3 in (66). The new mechanism generates much more NO than Fenimore‘s mechanism in the early stages of combustion, as seen in Fig. 1, because of the high [CH3] and [CH3O]. In this region, the reverse of reaction (67) is an additional source of N2O, without which the production of HCN via reactions initiated by N2O reacting with CH3 in (65) would die out. However, even in those parts of a flame where CH3O is almost fully consumed, the contribution of the proposed mechanism is higher than of Fenimore’s reactions because of more HCN being produced earlier.
Fig. 6. Comparison of [N2O] and [N2] measured by Vandooren et al. [23] along a CH4 ⫹ N2O ⫹ Ar flame and modeled by the proposed mechanism.
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Fig. 7. Comparison of [CH4], [CH3], and [NO] measured by Vandooren et al. [23] along CH4 ⫹ N2O ⫹ Ar flame and modeled by the mechanism proposed.
A crucial test of the proposed mechanism was the modeling of the experiment by Vandoren et al. [23]. This is presented in Figs. 6 – 8. Only in the case of NCO and N2 can some deficiency of the modeling be observed. It must be stated here that axial diffusion played an important role in this low pressure flame, without which mainly the position of the calculated curves would be delayed in comparison with the experiments. As can be seen HCN is formed in this flame according to the proposed mechanism and not by Fenimore’s reactions. The values of the rate coefficients for reactions (65) and (67) are crucial for modeling CH3. Without these reactions the maximum in [CH3] would be several times higher than measured by Vandooren et al. [23]. Figure 8 shows that the rate constants from Table I reproduced the measured maximum in [CH3].
319
Fig. 8. Comparison of the [HCN], [NCO], and [HNCO] measured by Vandooren et al. [23] along the flame of CH4 ⫹ N2O ⫹ Ar and modeled by the mechanism proposed.
Simplified mechanism forming NO in flames of CH4 ⴙ air The detailed combustion mechanism [10,27] permits a full analysis of a flames structure and also a prediction of the concentrations both the radicals and the stable species in the flame. However, longer times are require for computing. For this reason the application is limited to only simple one-dimensional laminar flames or perfectly stirred reactors. For more a complex geometry in e.g. industrial diffusion flames, the calculations become very expensive. This economical problem may be solved using simplified mechanisms for both combustion and the formation of pollutants. Assuming that [NO] is far from equilibrium, as is often the case in practical applications, the following simplified formula may be used for the rate of forming NO:
d关NO兴/dt ⫽ 2k1f关N2兴关O兴 ⫹ 2k4f关N2O兴关O兴 ⫹ 2k6b关N2O兴关H兴⫹2k10f 关NNH兴关O兴 ⫹ 2k56f关N2兴关CH兴 ⫹ ⫹ 2k65f 关N2O兴关CH3兴, where kif and kib, respectively, denote the rate constants of the forwards and backwards steps of reaction i. The terms in Eq. (I) represent the rate of NO formation from: Zeldovich’s reactions (1) and (2) with the assumption of a steady state for atomic nitrogen [1]; from N2O via (4) and the reverse of (6); the fourth term represents NO from NNH; the fifth term is Fenimore’s mechanism and finally are the reactions initiated by N2O reacting with CH3, i.e. the new mechanism. It was assumed that NH,
(I) N, and HCN are all fully converted to NO [28,29]. In and near the reaction zone, where large concentration of the free radicals H and OH are present, the N2O formed in the reverse of reactions (3) and (67) may be destructed not only by reactions (4), (5) and (65), but also by reactions (6), (7) and (8). Thus [N2O] within a flame quickly reaches a semisteady state at ⬃1 ppm. In accordance with the steady state assumption, [N2O] can be calculated from:
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关N2O兴 ⫽ 兵k3b关N2兴关O兴关M兴 ⫹ k67b关CH3O兴关N2兴其/兵k3f关M兴 ⫹ 共k4f ⫹ k5f兲关O兴 ⫹ (k6b ⫹ k7f兲关H兴 ⫹ ⫹ k8f关OH兴 ⫹ 共k65f ⫹ k67f兲关CH3兴} Glarborg et al. [30], on the basis of experiments conducted at temperatures below 1400 K, suggested that reaction (8) may play a minor role destroying N2O. They found that its rate constant is ⬃10 times lower than the values recommended by Miller and Bowman [10] and ⬃60 times lower than the values recommended by Tsang and Herron [31]. A similar conclusion has been reported by Sausa et al. [32]. Despite this, reaction (8) was included in Eq. (II). In accordance with Bozzelli and Dean’s [13] suggestion, [NNH] can be calculated assuming that the fast reaction (9) is equilibrated, so [NNH] ⫽ [N2] [H]/K9, where K9 is the equilibrium constant of (9). Application of the above equations requires a knowledge of [O], [H], [OH], [CH], [CH3], and [CH3O]. Equation (I) is able to predict well the rates of formation of NO in a flame, if the concentrations of the main radicals are properly established. To simplify calculating the concentrations of the radicals, many ways based on partial equilibrium and the steady state assumption were proposed [22,29,33– 36]. The difference between the detailed mechanism in Table I and the above simplified formulae, which require the rate constants for only 12 reactions: (1), (3–10), (56), (65) and (67), depends on the accuracy of these simplified calculations, of which the least accurate are for [CH3] and [CH3O]. The presence of methane within a flame creates problems for calculating [O], [H] and [OH] using partial equilibrium in both the systems of H2/O2 and CO/H2/O2 as already addressed [29] and demonstrated in Figs. 9 and 10 for of distribution of [O] and [OH] in a premixed, moderately lean (air ⫽ 1.1 x stoichiometric amount) flame of CH4 ⫹ air at constant temperature of 1800 K. The results obtained from the classical partial equilibrium assumption of H2/O2 [22,33] agree very well with the ones calculated using the detailed combustion mechanism, but only in the region where CH4 is completely consumed. Calculations at lower temperatures of 1700 and 1600 K gave similar results. Some of the proposed reduced mechanisms of combustion allow one to compute [H] [34,35,37]. In this case, if the local [H] in the flame is known, [O] and [OH] can only be obtained from assuming partial equilibrium of H2/O2 [35]. In contrast with the classical H2/O2 partial equilibrium, the [O] and [OH] obtained by this approach agree very well with those calculated from the detailed combustion mechanism. Figure 11 compares the rate of formation of NO and also [NO] modeled in the same CH4 ⫹ air flame (air ⫽ 1.1 x stoichiometric, T ⫽ 1800 K) as derived
(II)
Fig. 9. Modeled [O] and [CH4] in a flame of CH4 ⫹ air (air ⫽ 1.1 ⫻ stoichiometric, T ⫽ 1800 K): (——) [O] from detailed combustion mechanism, (— —) [O] from classical H2/O2 partial equilibrium, (- - -) [O] from H2/O2 partial equilibrium with [H] obtained from detailed combustion mechanism.
from the detailed computations for both combustion [10] and the production of NO, with the mechanism in Table I and secondly as derived from Eq. (I) using the mentioned two above-ways of assuming partial equilibrium of H2/O2. For hydrocarbon radicals in Eq. (I), expressions based on the steady state assumption were used: these were a slightly modified Eq. (3) in [36] for [CH3] and Eq. (44) in [29] for [CH]. The
Fig. 10. Modeled [OH] and [CH4] in the same flame as Fig. 9: (——) [OH] from detailed combustion mechanism, (— —) [OH] from classical H2/O2 partial equilibrium, (- - -) [OH] from H2/O2 partial equilibrium with [H] obtained from detailed combustion mechanism.
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321
assumed to be fast. Further investigations of the proposed reactions are necessary in order to improve the modeling of NO in hydrocarbon flames. The simplified mechanism (Eq.I) for the formation of NO, together with the assumption of partial equilibrium for [NNH], [O] and [OH] and steady state for [N2O], [CH], [CH3] and [CH3O] was found to be a good approximation for modeling [NO] in industrial flames of CH4 ⫹ air.
References Fig. 11. NO formation rate and [NO] in the same flame as Fig. 9 calculated from the detailed combustion and NO mechanisms (——) and also the simplified NO mechanism (Eqs.I-II) with [CH], [CH3] and [CH3O] from the steady state assumption and: (— —) [H], [O], [OH] from the classical H2/O2 partial equilibrium, (- - -) [O], [OH] from H2/O2 partial equilibrium and [H] from the detailed combustion mechanism.
steady state assumption was used also for calculating [CH3O]. As is seen from Fig. 11 the rates of formation of NO obtained with the classical H2/O2 partial equilibrium are very high, probably as a consequence of [O] being too large. However, in the case when only [O] and [OH] were evaluated using partial equilibrium of H2/O2 with [H] from the detailed combustion mechanism, the rates of formation of NO obtained from Eq. (I) are very close to those calculated from the detailed mechanism for combustion and NO. At short times, the simplified mechanism overpredicts [NO] due to the fact that formation and conversion of HCN into NO are assumed to proceed simultaneously, while in the detailed mechanism a considerable delay time can be noticed. However, the final [NO] is smaller because of this simplification.
Conclusions The proposed reactions of N2O and NNH with CH3 radicals provide an additional mechanism forming NO in hydrocarbon flames. The role of this mechanism depends on the rate constants of the proposed reactions, which at present could only be estimated. Nevertheless, the rate constants of the new reactions proposed in Table I result from an analysis of experimental measurements. Introduction of new reactions always creates problems, but all the experiments analyzed [15,21– 23] could only be described if the new reaction N2O ⫹ CH3 3 CH2NH ⫹ NO and the previously studied reaction N2O ⫹ CH3 7 CH3O ⫹ N2 [24,25] were
[1] Y.B. Zeldovich, Acta Physicochemica USSR, 21 (1946) 577. [2] C.P. Fenimore, Thirteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1971, p. 373. [3] G.G. de Soete, Fifteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1975, p. 1093. [4] A.D. Al-Fawaz, L. M.Dearden, J.T. Hedley, H. Missaghi, M. Pourkashanian, A. Williams, L.T. Yap, Twenty-Fifth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1994, p. 1027. [5] A.N. Hayhurst, I.M. Vince, Prog. Energy Combust. Sci. 6 (1980) 35. [6] C.P. Fenimore, H.A. Fraenkel, Eighteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1981, p. 143. [7] J. Luque, G.P. Smith, D.R. Crosley, Twenty-Sixth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1996, p. 959. [8] J. Wolfrum, Chemie-Ingeneur-Technik, 44 (1972) 656. [9] P.C. Malte, D.T, Pratt, Combust. Sci. Technol. 9 (1974) 221. [10] J. A. Miller, C.T. Bowman, Prog. Energy Comb. Sci. 15 (1989) 287. [11] J. Tomeczek, B. Gradon´ , Combust. Sci. Technol. 125 (1997) 159. [12] J. Tomeczek, B. Gradon´ , Fifth International Conference on Technologies and Combustion for a Clean Environment, Lisbon-Portugal, 1999, p. 129. [13] J.W. Bozzelli, A.M. Dean, Int. J. Chem. Kin. 27 (1995) 1097. [14] J. Blauwens, B. Smets, J. Peeters, Sixteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1977, p. 1055. [15] A.N. Hayhurst, E.M. Hutchinson, Combust. Flame 114 (1998) 274. [16] A.A. Konnov, G. Colson, J. de Ruyck, Combust. Flame 121 (2000) 548. [17] A.A. Konnov, and J. de Ruyck, Combust. Flame 125 (2001) 1258. [18] J. Tomeczek, B. Gradon´ , Sixth International Conference on Technologies and Combustion for a Clean Environment, Porto-Portugal, 2001, p. 513.
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[19] S.K. Vidyarthi, C. Willis, R.A. Back, J. Phys. Chem. 80 (1976) 559. [20] Burcat, A., ftp://ftp.technion.ac.il/pub/supported/aetdd/ thermodynamics (date accessed, 2002). [21] P.C. Malte, S.C. Schmidt, D.T. Pratt, Sixteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1977, p. 145. [22] A.F. Sarofim, G. Pohl, Fourteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1973, p. 739. [23] J. Vandooren, M.C. Branch, P.J. Van Tiggelen, Combust. Flame 90 (1992) 247. [24] J.W. Falconer, D.E. Hoare, R. Overend, Combust. Flame 21 (1973) 339. [25] A.A. Borisov, W.M. Zamanskij, K. Potmishil, G.I. Skatshkov, W.A. Foteenkov, Kinetika i Kataliz 18 (1977) 307. [26] J.W. Bozzelli, A.Y. Chang, A.M. Dean, Twenty-Fifth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1994, p. 965. [27] P. Glarborg, J.A. Miller, R.J. Kee, Combust. Flame 65 (1986) 177. [28] R.C. Steele, P.C. Malte, D.G. Nicol, J.C. Kramlich, Combust. Flame 100 (1995) 440. [29] A.A. Konnov, NO Formation Rates in Natural Gas Combustion. Fourth International Conference on Technologies and Combustion for a Clean Environment, Lisbon-Portugal, 1997, p. 9. [30] P. Glarborg, J.E. Johnsson, K. Dam-Johansen, Combust. Flame 99 (1994) 523.
[31] W. Tsang, J.T. Herron, J. Phys. Chem. Ref. Data 20 (1991) 609. [32] R.C. Sausa, W.R. Anderson, D.C. Dayton, C.M. Faust, S.L. Howard, Combust. Flame 94 (1993) 407. [33] D. Thompson, T.D. Brown, J.M. Beer, Fourteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1973, p. 787. [34] G. Paczko, P.M. Lefdal, N. Peters, Twenty-First Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1986, p. 739. [35] N. Peters, R.J. Kee, Combust. Flame 68 (1987) 17. [36] P. Glarborg, N.I. Lilleheie, S. Byggstoyl, B.F. Magnussen, P. Kilpinen, M. Hupa, Twenty-Fourth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1992, p. 889. [37] J.M. Card, W.T. Ashurst, F.A. Williams, Combust. Flame 97 (1994) 48. [38] J.D. Mertens, A.Y. Chang, R.K. Hanson, C.T. Bowman, Int. J. Chem. Kin. 23 (1991) 173. [39] K.A. Miller, C.F. Melius, Twenty-Fourth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1992, p. 719. [40] J. Vandooren, P.J. Van Tiggelen, J.F. Pauwels, Combust. Flame 109 (1997) 647. [41] B. Gradon´ , J. Tomeczek, Combust. Flame 126 (2001) 1856. [42] P. Glarborg, P.G. Kristensen, S.H. Jensen, K. DamJohansen, Combust. Flame 98 (1994) 241. [43] http://www.me.berkeley.edu/grimech/version30/files30/ grimech30.da (date accessed, 1999).
The role of N2O and NNH in the formation of NO via HCN in hydrocarbon flames J. Tomeczek*, B. Gradon´ Department of Process Energy, Silesian Technical University, ul.Krasin´skiego 8, 40-019 Katowice, Poland Received 14 April 2002; received in revised form 2 January 2003; accepted 14 January 2003
Abstract A new way of forming HCN in flames via N2O and NNH reacting with CHi radicals is proposed and tested for rich and lean gaseous premixed flames of CH4 and air and also of CH4, N2O and Ar. This new route is thermodynamically more probable than Fenimore’s direct reaction of N2 with CHi radicals. In fact, it is shown that the new mechanism is more important than Fenimore’s reaction in both rich and lean flames. Rate constants of the new reactions forming NO have been suggested on the basis of numerical modeling. It has been shown that the formation of NO through HCN is most effective as the result of reactions initiated by N2O ⫹ CH3 3 CH2NH ⫹ NO, followed by CH2NH ⫹ H 3 H2CN ⫹ H2 and CH2NH ⫹ O 3 H2CN ⫹ OH. In flames of CH4 and air, a substantial source of N2O comes from the reverse of the reaction N2O ⫹ CH3 7 CH3O ⫹ N2 in the reaction zone. A formula based on the steady state assumption and partial equilibrium limits the number of nitrogen conversion reactions to only 12; this was tested using a premixed flame of CH4 and air. © 2003 The Combustion Institute. All rights reserved. Keywords: HCN; N2O; NNH and NO formation
Introduction Modeling [NO] in gaseous hydrocarbon flames still creates fundamental difficulties. There are two distinct problems: the details of the mechanism of chemical reactions producing NO, together with values of their rate constants; the concentration of radicals participating in this mechanism, mainly [O], [H] and [CHi]. Zeldovich [1] proposed a formation mechanism of two reactions producing thermal NO in: N2 ⫹ O 3 NO ⫹ N,
(1)
N ⫹ O2 3 NO ⫹ O,
(2)
and calculated [O] assuming that the thermal disso* Corresponding Author. Tel./Fax: ⫹48-32-603-4286. E-mail address: [email protected] (J. Tomeczek).
ciation of O2 in O2⫹ M N O ⫹ O ⫹ M was at equilibrium. This assumption enabled him to find the rate constants of the thermal NO mechanism. This approach was fully justified, because the flame temperature in Zeldovich’s experiments was above 2000°C. In many flames, however, [O] is much higher than [O]eq calculated from thermal equilibrium, but even using the real [O], it is not possible to explain the observed [NO] in hydrocarbon and other flames using Zeldovich’s reactions (1) and (2). Fenimore [2] accordingly proposed the so-called prompt reaction between molecular nitrogen N2 and CHi radicals in hydrocarbon flames. For the Fenimore’s mechanism, overall rate equations were proposed [3,4], which often were based on subtracting from the observed rate of NO the rate of Zeldovich’s reactions (1) and (2) calculated using [O]eq. Such equations include any reaction taking place within the flame and not only Fenimore’s reaction, which
0010-2180/03/$ – see front matter © 2003 The Combustion Institute. All rights reserved. doi:10.1016/S0010-2180(03)00013-0
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was extensively studied [5]. Indeed, its contribution was postulated to be important in cooler, rich hydrocarbon flames, whereas in hotter, lean flames, the amount of prompt NO fell abruptly. Fenimore and Fraenkel [6] later pointed out that the role of the N2 ⫹ CHi had been probably overvalued. Luque et al. [7] found that the rate of Fenimore’s principal reaction N2 ⫹ CH 3 HCN ⫹ N
(56)
should be about twice as fast than traditionally assumed to reproduce the measured [NO] in flames of methane ⫹ air. Wolfrum [8] and Malte and Pratt [9] proposed the nitrous oxide mechanism based on three reactions: N2O ⫹ M7N2 ⫹ O ⫹ M,
(3)
N2O ⫹ O 3 NO ⫹ NO,
(4)
N 2 O ⫹ O 3 N2 ⫹ O2 .
(5)
However, the recommended rate constants of these reactions [10] limited the contribution of this mechanism to only ⬃12% of the rate of formation of NO. Tomeczek and Gradon´ [11] found that the thermal mechanism is much faster than described by the Zeldovich reactions with the recommended rate constants. They explained it by introducing the term “extended thermal mechanism” including both Zeldovich’s reactions (1 and 2) and the nitrous oxide reactions (3 to 5), for which the rate constants of the controlling reactions were proposed and substantiated with experimental evidence from the literature. The extended thermal mechanism can reproduce the observed [NO] in lean CH4 ⫹ air flames and in rich H2 ⫹ CO ⫹ N2O ⫹ Ar flames [12] if proper [O] are used. However, for rich flames of CH4 ⫹ air, the calculated [NO] are still too small; this suggests that further reactions operate in these conditions. Bozzelli and Dean [13] introduced the NNH mechanism composed of the two main reactions: NNH ⫹ M7N2 ⫹ H ⫹ M, NNH ⫹ O 3 NH ⫹ NO,
(9) (10)
for which they proposed the kinetic constants. They came also to the conclusion that within a flame the first reaction is close to equilibrium. The aim of this research was to tackle the modeling of [NO] in rich and near-stoichiometric flames of CH4 ⫹ air using available experimental data for gaseous fuels. However, rich flames do not occur often. Nevertheless, in practical diffusion flames, there are large volumes containing rich mixtures where most of the NO can be formed. The difficulty has been that any additional mechanism must be able
to increase [NO] in rich conditions, but at the same time to give less NO when the flame’s composition shifts toward lean conditions. It has been found that these difficulties can be solved using the new proposed reactions of CHi radicals with N2O and NNH. These, together with the extended thermal mechanism and the modified NNH mechanism, predict very well [NO] in hydrocarbon flames. A simplified version of the mechanism is proposed also for modeling large practical flames and is tested with a flame of CH4 ⫹ air.
Proposed mechanism for the formation of HCN Following Fenimore [2], a direct attack of a hydrocarbon radical on molecular N2 is traditionally considered as the first step on the road to NO via HCN as an intermediate [5]. However, the reactions of CHi with N2 are endothermic with positive values of the change in Gibbs free energy ⌬Go, thus giving very small equilibrium constants for these reactions, so the contribution to [NO] must be insignificant in most flame conditions. Blauwens et al. [14] on the basis of mass spectrometric measurements found that it is impossible to distinguish between reactions (56) and (57). Miller and Bowman [10] concluded that only Fenimore’s principal reaction (56) cannot be ruled out; however, they also included reaction (57) in their system of reactions for producing NO. The shortcomings of Fenimore’s mechanism are possibly overcome by the NNH mechanism; however, the results must be taken cautiously. Hayhurst and Hutchinson [15] could not find any evidence for the production of NO via N2O, but confirmed that significant amounts of NO are formed via NNH through reaction (10) in rich flames of H2 ⫹ O2 ⫹ N2. This has been tested by Konnov et al. [16]. As a matter of fact, the difference between the experimental [NO] and those modeled by Miller and Bowman’s [10] mechanism were taken [15] as a contribution to reaction (10), with subsequent total conversion of NH into NO. It is necessary to emphasize that most of these difficulties can be explained using the extended thermal mechanism proposed by Tomeczek and Gradon´ [11], where N2O plays a very important role. What is more, there is direct experimental evidence, gathered from the literature by Tomeczek and Gradon´ [11], that the rate of reaction (4) is much faster than traditionally assumed, while for the reaction (10) the rate constant was not measured directly. Even in a flow of air without fuel through a high temperature reactor, the measured [NO] at the outflow will be from 3 to 5 times higher [11] than predicted by the Zeldovich mechanism with Miller and Bowman’s [10] rate constants. The NNH mech-
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
anism with the available rate constants together with the extended thermal mechanism [11] considerably overpredicts [NO] in lean CH4 ⫹ air flames [12], which suggests that the rate coefficient of reaction (10) proposed by Bozzelli and Dean [13] and Hayhurst and Hutchinson [15], recently corrected by Konnov and de Ruyck [17], is too high. The presence of HCN in flames with a hydrocarbon as fuel is an experimental fact, but in addition to (56) and (57) it can also be formed [18] exothermically by hydrocarbon (CHi) fragments reacting with N2O and NNH. The elementary steps of these reactions are discussed below.
flames should be expected for (66). Nevertheless, it has been included in the final mechanism, but with a low rate coefficient. The interesting feature of reaction (67) is that in some flame regions it can produce N2O as a result of high [N2] and [CH3O]. This reaction is exothermic; nevertheless, in regions of high [CH3O], thermodynamics (⌬G0 ⫽ ⫺218.3 kJ/mol) favors its backwards step. The newly proposed reverse of reaction (67) increases [N2O] so much that N2O can become a dominant source for HCN. The subsequent reactions most probably leading to the formation of HCN are: CH2NH ⫹ H 3 H2CN ⫹ H2, ⌬H ⫽ ⫺ 54.6 kJ/mol, 0
Discussion
313
(68) (68)
CH2NH ⫹ O 3 H2CN ⫹ OH, The state of the art leads one to expect that the role of NNH in the formation of HCN should be much more important than that of N2O, because of the high rate of formation [13] of NNH. The introduction of new reactions consuming NNH should not influence its concentration due to [NNH] being close to equilibrium in flames [13]. On the other hand, the reactions of CH3 radicals are most likely to produce enough HCN in flames, because [CH] and [CH2] are lower than [CH3]. The reaction of NNH with CH3 can proceed to a known products in: NNH ⫹ CH3 3 N2 ⫹ CH4, ⌬H ⫽ ⫺471.3 kJ/mol, 0
(64)
N2O ⫹ CH3 3 CH2NH ⫹ NO, (65)
NNH ⫹ CH3 3 CH2N2 ⫹ H2, ⌬H0 ⫽ ⫺ 110.3 kJ/mol.
(66)
However the reactants of (65) can also undergo the parallel reaction: N2O ⫹ CH37CH3O ⫹ N2, ⌬H0 ⫽ ⫺ 212.3 kJ/mol.
(69)
CH2N2 ⫹ O 3 H2CN ⫹ NO, ⌬H0 ⫽ ⫺ 196.9 kJ/mol,
(71)
H2CN ⫹ H 3 HCN ⫹ H2, ⌬H0 ⫽ ⫺ 334.7 kJ/mol,
(72)
H2CN ⫹ O 3 HCN ⫹ OH, ⌬H0 ⫽ ⫺ 326.5 kJ/mol,
(73)
CH2N2 ⫹ O 3 HCNN ⫹ OH,
which does not yield NO. It is also possible that the same reactants initially form the stable molecule N2CH4 [19] in: NNH ⫹ CH3 ⫹ M 7 N2CH4 ⫹ M, which must have an unrealistically high rate coefficient in order to contribute to the formation of HCN. An alternative way to HCN is through the known molecules CH2N2 and CH2NH [20], which can be formed in:
⌬H0 ⫽ ⫺53.3 kJ/mol.
⌬H0 ⫽ ⫺ 46.4 kJ/mol,
(67)
The four-centered reaction (66) is similar in character to Fenimore’s second reaction (57); thus, a minor role in the formation of NO in premixed hydrocarbon
⌬H0 ⫽ ⫺ 33.9 kJ/mol.
(74)
Thus, the H2CN formed in reactions (68) and (69) may be converted into HCN in reactions (72), (73), and (76), while the HCNN formed in reaction (74) may give HCN in reaction (100). The rate constants of these reactions can only be suggested, because at present no reliable experimental meassurements are available. Thus, a numerical, modeling way has been chosen to adjust the rate constants. The published chemical compositions of the flames of CH4 ⫹ air [15,21,22] do allow for the rate coefficients to be determined, but only with an activation energy of zero, which is controversial. This problem of the activation energy could be partially solved using the detailed measurements of the flame of CH4 ⫹ N2O ⫹ Ar by Vandooren et al. [23], because of a broad range and variation of stable species and radicals in this flame. However, a scheme of only the above reactions was unable to describe the composition of this flame, which was very difficult to model. Particular problems were to compute the concentrations of HNCO, HCN, and NO. The known reactions involving HNCO (107-122) could not solve the problem; thus the following new reactions facilitated the modeling:
314
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N2O ⫹ HCN 3 HNCO ⫹ N2, ⌬H0 ⫽ ⫺330.7 kJ/mol,
(104)
HNCO ⫹ NO 3 NNH ⫹ CO2, ⌬H0 ⫽ ⫺117.2 kJ/mol,
(105)
HNCO ⫹ NH27NCO ⫹ NH3, ⌬H0 ⫽ ⫹12.0 kJ/mol.
(106)
When formulating rate coefficients, the pre-exponential factor was kept a safe margin below the upper limit imposed by kinetic theory. For all the proposed reactions, the suggested pre-exponential factors and activation energies are given in Table I. The crucial reaction forming HCN in the proposed mechanism is (65). However, without the reverse of reaction (67) as the source of N2O, (65) is not effective. Both these reactions must be fast, otherwise the measured profiles of [N2O] and [CH3] along Vandooren et al.’s [23] flame cannot be described. As far as reaction (65) is concerned, there is no previous work for comparison; however, for reaction (67) there are two reports suggesting that at 873 K [24,25], this reaction is slow with a rate coefficient of ⬃14 m3/(mol s). These results were obtained from simplified modeling involving only 11 reactions; in fact, many of the species measured in large quantities by Vandooren et al. [23] were not included. Using the rate coefficient of reaction (67) assumed by Borisov et al. [25], the experiment by Vandooren et al. [23] could not be described. Our finding was that the rate coefficients suggested in Table I provided the only solution to match experimental evidence. Otherwise, Miller and Bowman’s [10] combustion mechanism was assumed. Numerical simulations for rich CH4 ⫹ air flames have shown that HCN is formed in the new reactions earlier than in Fenimore’s principal reaction (56), because CH3 precedes CH in such a flame. The contribution of the proposed reactions was tested for a rich mixture (air ⫽ 0.9 stoichiometric amount) of CH4 ⫹ air of constant temperature 1800 K. The results are presented in Fig. 1. The new mechanism generates more HCN than Fenimore’s reactions, for which the [HCN] profile is very narrow. However, the maximum in [HCN] for both mechanisms is observed in similar positions, so [NO] predicted by the new mechanism is also higher than described by Fenimore’s mechanism. The CH4 flames analyzed by Malte et al. [21], Sarofim and Pohl [22], and by Hayhurst and Hutchinson [15] were chosen to verify the modeling of hydrocarbon flames. Altogether, four different mechanisms were analyzed: three (Miller and Bowman [10], extended thermal mechanism (1 to 5); [11]
extended thermal mechanism (1 to 5) plus reaction (6) from Bozzelli et al. [23]) reflected current knowledge, but the fourth new one comprised reactions (1 to 122) from Table I. Figure 2 presents the results for three cases of modeling using available mechanisms in the literature for the experimental measurements of Malte et al. [21]. In the modeling, the flame temperature was assumed to be the mean of the values measured by a thermocouple and by the OH rotational method. Miller and Bowman’s [10] mechanism produces [NO] much lower than the experimental values for both the lean and rich flames. The extended thermal mechanism is also not satisfactory, even after including reaction (6). For very lean flames, the extended thermal mechanism is able to account for over 50% of the observed [NO]. The [NO] generated in a slightly rich flame of CH4 ⫹ air as investigated by Sarofim and Pohl [22] are presented in Fig. 3, together with the modeling results for the three mechanisms representing the state of art. Again, the extended thermal mechanism with the additional reaction (6) produces reasonable agreement with the experimental points. For the very rich CH4 ⫹ air flame of Hayhurst and Hutchinson [15], the extended thermal mechanism with reaction (6) also gives the smallest divergence from the experimental points as seen in Fig. 4. However, still over 20% of the [NO] is not predicted. None of the three mechanisms analyzed using known rate constants can describe satisfactorily the observed [NO]. This suggests that some important reactions have been neglected. To be objective, the consequences of modifying of the Fenimore’s reactions by enhancing of their rate constants were examined. In this way it was possible to improve the modeling in rich flames, but unfortunately in flames near to stoichiometric and in lean ones the modeled of [NO] were too high. Assessment of reaction (10) has to take into account the additional contribution of the new reactions of N2O and NNH with CHi radicals. A proper value of the (10) reaction rate constant can be found only using hydrogen flames to exclude the contribution of HCN. The difference between the experimental [NO] values and those calculated by the extended thermal mechanism with the rate constants by Tomeczek and Gradon´ [11] has been taken as the contribution of reaction (10) in hydrogen flames; the evaluated rate constants for reaction (10) are given in Table I, and are smaller than those proposed by Bozzelli and Dean [13], because they used Miller and Bowman’s [10] rate constants for the thermal mechanism. The final consequence for modeling NO is demonstrated successfully in Fig. 5 for a very rich H2 ⫹ air flame investigated by Hayhurst and Hutchinson [15]. Three
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
315
Table I Rate constants of nitrogen conversion reactions, kf ⫽ k0 Tn exp(⫺E/RT) No. 1 2 3 Enhanced 4 5 6 7 8 9 Enhanced 10 11 12 13 14 15 16 17 18 19 20 Enhanced 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 Enhanced 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
2.72 ⫻ 108 N2⫹O7NO⫹N 0.00 6.40 ⫻ 103 N⫹O27NO⫹O 1.00 3.20 ⫻ 108 N2O⫹M7N2⫹O⫹M 1.00 third-body efficiencies: O2 (1.4), CO2 (3.0), N2 (1.7), H2O (12.0), other (1.0) N2O⫹O7NO⫹NO 6.21 ⫻ 108 0.00 N2O⫹O7N2⫹O2 6.21 ⫻ 108 0.00 2.50 ⫻ 1010 NH⫹NO7N2O⫹H ⫺1.03 2.20 ⫻ 108 N2O⫹H7N2⫹OH 0.00 N2O⫹OH7N2⫹HO2 1.64 ⫻ 105 0.00 1.30 ⫻ 108 ⫺0.11 NNH⫹M7N2⫹H⫹M third-body efficiencies: all species (1.0) NNH⫹O7NO⫹NH 3.30 ⫻ 107 ⫺0.23 2.90 ⫻ 105 NNH⫹O27N2O⫹OH ⫺0.34 1.40 ⫻ 108 NNH⫹O7N2O⫹H ⫺0.40 NNH⫹H7N2⫹H2 1.00 ⫻ 105 0.00 2.40 ⫻ 1016 NNH⫹OH7N2⫹H2O ⫺2.88 1.70 ⫻ 1010 NNH⫹O7N2⫹OH ⫺1.23 NNH⫹O27N2⫹HO2 1.20 ⫻ 106 ⫺0.34 5.00 ⫻ 107 NNH⫹NO7N2⫹HNO 0.00 NNH⫹NH7N2⫹NH2 5.00 ⫻ 107 0.00 NNH⫹NH27N2⫹NH3 5.00 ⫻ 107 0.00 1.10 ⫻ 1010 NO2⫹M7NO⫹O⫹M 0.00 third-body efficiencies: all species (1.0) 8.43 ⫻ 107 NO2⫹H7NO⫹OH 0.00 NO2⫹O7NO⫹O2 3.91 ⫻ 106 0.00 2.11 ⫻ 106 NO⫹HO27NO2⫹OH 0.00 7.60 ⫻ 104 NH⫹O27NO⫹OH 0.00 NH⫹O7NO⫹H 5.50 ⫻ 107 0.00 5.00 ⫻ 105 NH⫹OH7N⫹H2O 0.50 3.00 ⫻ 107 NH⫹N7N2⫹H 0.00 NH⫹O7N⫹OH 3.72 ⫻ 107 0.00 2.16 ⫻ 107 NH⫹NO7N2⫹OH ⫺0.23 1.60 ⫻ 108 N⫹H27NH⫹H 0.00 5.10 ⫻ 107 NH⫹NH7N2⫹H⫹H 0.00 NH⫹OH7NO⫹H2 2.00 ⫻ 107 0.00 3.89 ⫻ 107 NH⫹O27HNO⫹O 0.00 NH⫹OH7HNO7H 2.00 ⫻ 107 0.00 HNO⫹M7NO⫹H⫹M 1.50 ⫻ 1010 0.00 third-body efficiencies: H2O (10.0), O2 (2.0), N2 (2.0), H2 (2.0), other (1.0) HNO⫹OH7NO⫹H2O 4.80 ⫻ 107 0.00 HNO⫹O7NO⫹OH 3.61 ⫻ 107 0.00 1.81 ⫻ 107 HNO⫹H7H2⫹NO 0.00 3.95 ⫻ 106 HNO⫹HNO7N2O⫹H2O 0.00 2.00 ⫻ 106 HNO⫹NO7N2O⫹OH 0.00 N⫹OH7NO⫹H 3.80 ⫻ 107 0.00 7.00 ⫻ 106 NH2⫹O7OH⫹NH 0.00 6.63 ⫻ 108 NH2⫹O7HNO⫹H ⫺0.50 4.00 2.00 NH2⫹OH7NH⫹H2O NH2⫹NO7NNH⫹OH 6.40 ⫻ 109 ⫺1.25 6.20 ⫻ 1012 NH2⫹NO7N2⫹H2O ⫺1.25 7.20 ⫻ 107 NH2⫹N7N2⫹H⫹H 0.00 4.50 ⫻ 106 NH2⫹O27HNO⫹OH 0.00 NH2⫹H7NH⫹H2 6.92 ⫻ 107 0.00 2.00 ⫻ 107 NH2⫹HNO7NH3⫹NO 0.00 1.40 ⫻ 1010 NH3⫹M7NH2⫹H⫹M 0.00
E kJ mol⫺1
References
318.59 26.19 232.13
[11] [10] [30]
108.81 108.81 3.49 70.10 41.57 20.84
[11] [11] [23] [23] [30] [13]
⫺4.23 0.62 1.99 0.00 10.23 2.08 0.62 0.00 0.00 0.00 274.36
[18] [13] [13] [13] [13] [13] [13] [10] [10] [10] [10]
0.00 1.00 ⫺2.01 6.41 0.00 8.37 0.00 0.00 0.00 105.26 0.00 0.00 74.88 0.00 203.82
[31] [31] [10] [10] [38] [10] [10] [38] [39] [32] [32] [40] [38] [10] [10]
4.14 0.00 4.16 20.93 108.86 0.00 0.00 0.00 4.19 0.00 0.00 0.00 104.67 15.28 4.19 379.34
[31] [31] [31] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [27] continued
316
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
Table I Continued No.
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
Enhanced third-body efficiencies: all species (1.0) 5.00 ⫻ 1010 NH3⫹M7NH⫹H2⫹M 0.00 Enhanced third-body efficiencies: all species (1.0) 53 NH3⫹H7NH2⫹H2 0.64 2.39 2.04 2.04 54 NH3⫹OH7NH2⫹H2O 55 NH3⫹O7NH2⫹OH 2.10 ⫻ 107 0.00 3.00 ⫻ 105 56 N2⫹CH7HCN⫹N 0.00 1.00 ⫻ 107 57 N2⫹CH27HCN⫹NH 0.00 5.00 ⫻ 107 58 N⫹CH37HCN⫹H⫹H 0.00 0.00 59 NO⫹CH7HCN⫹O 1.10 ⫻ 108 1.00 ⫻ 105 60 NO⫹CH37HCN⫹H2O 0.00 2.00 ⫻ 107 61 NO⫹CH27HCN⫹OH 0.00 8.00 ⫻ 107 62 N2O⫹CH7HCN⫹NO 0.00 63 NNH⫹CH7HCN⫹NH 8.00 ⫻ 107 0.00 64 NNH⫹CH37N2⫹CH4 2.50 ⫻ 107 0.00 6.73 ⫻ 104 65 N2O⫹CH37CH2NH⫹NO 2.00 66 NNH⫹CH37CH2N2⫹H2 1.00 ⫻ 107 0.00 67 N2O⫹CH37CH3O⫹N2 1.90 ⫻ 104 1.75 68 CH2NH⫹H7H2CN⫹H2 8.32 ⫻ 103 1.25 1.62 ⫻ 104 69 CH2NH⫹O7H2CN⫹OH 1.25 4.00 ⫻ 107 70 CH2NH⫹O7CH2O⫹NH 0.00 8.00 ⫻ 107 71 CH2N2⫹O7H2CN⫹NO 0.00 72 H2CN⫹H7HCN⫹H2 5.00 ⫻ 107 0.00 5.00 ⫻ 107 73 H2CN⫹O7HCN⫹OH 0.00 2.00 ⫻ 107 74 CH2N2⫹O7HCNN⫹OH 0.00 75 H2CN⫹N7N2⫹CH2 6.00 ⫻ 107 0.00 3.30 ⫻ 101 76 HCN⫹H⫹M7H2CN⫹M 0.00 Enhanced third-body efficiencies: H2 (2.0), H2O (6.0), CH4 (2.0), CO (1.5), CO2 (2.0), C2H6 77 CH3⫹NO7H2CN⫹OH 1.00 ⫻ 106 0.00 6.10 ⫻ 108 78 CH3⫹N7H2CN⫹H ⫺0.31 79 HCN⫹O7NH⫹CO 3.45 ⫻ 10⫺3 2.64 80 HCN⫹O7NCO⫹H 1.38 ⫻ 10⫺2 2.64 7.83 ⫻ 10⫺10 81 HCN⫹OH7NH2⫹CO 4.00 82 NCO⫹H7NH⫹CO 5.00 ⫻ 107 0.00 83 NCO⫹O7NO⫹CO 2.00 ⫻ 107 0.00 2.00 ⫻ 107 84 NCO⫹N7N2⫹CO 0.00 85 NCO⫹M7N⫹CO⫹M 3.10 ⫻ 1010 ⫺0.50 Enhanced third-body efficiencies: N2 (1.5), O2 (1.5), H2O (18.6), other (1.0) 86 NCO⫹NO7N2O⫹CO 1.00 ⫻ 107 0.00 87 NCO⫹OH7NO⫹CO⫹H 1.00 ⫻ 107 0.00 1.90 ⫻ 105 88 N⫹CO27NO⫹CO 0.00 89 CH⫹N7CN⫹H 1.30 ⫻ 107 0.00 1.45 ⫻ 107 90 HCN⫹OH7CN⫹H2O 0.00 91 HCN⫹O7CN⫹OH 2.70 ⫻ 103 1.58 2.92 ⫻ 10⫺1 92 CN⫹H27HCN⫹H 2.45 93 CN⫹O7CO⫹N 1.80 ⫻ 107 0.00 5.60 ⫻ 106 94 CN⫹O27NCO⫹O 0.00 95 CN⫹OH7NCO⫹H 6.00 ⫻ 107 0.00 3.00 ⫻ 107 96 CN⫹NO27NCO⫹NO 0.00 97 CN⫹N2O7NCO⫹N2 1.00 ⫻ 107 0.00 3.10 0.15 98 CH⫹N2⫹M7HCNN⫹M Enhanced third-body efficiencies: H2 (2.0), H2O (6.0), CH4 (2.0), CO (1.5), CO2 (2.0), C2H6 99 HCNN⫹O7CO⫹H⫹N2 2.20 ⫻ 107 0.00 52
E kJ mol⫺1
References
340.00
[41]
42.59 2.37 37.68 56.94 309.83 0.00 0.00 62.80 0.00 0.00 0.00 0.00 127.00 0.00 80.00 10.00 20.00 20.00 0.00 0.00 0.00 0.00 1.67 0.00 (3.0), other (1.0) 91.07 1.21 20.85 20.85 16.75 0.00 0.00 0.00 201.00
[10] [10] [10] [10] [10] [27] [10] [10] [10] this work this work [43] this work this work this work this work this work this work this work this work this work this work [43] [43]
⫺1.63 0.00 14.24 0.00 45.76 111.37 9.37 0.00 0.00 0.00 0.00 0.00 0.00 (3.0), other (1.0) 0.00
[10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [10] [43]
[43] [43] [10] [10] [10] [10] [10] [10] [42]
[43] continued
J. Tomeczek, B. Gradon´ / Combustion and Flame 133 (2003) 311–322
317
Table I Continued No.
Reaction
k0 m3K⫺n mol⫺1 s⫺1
n ⫺
E kJ mol⫺1
References
100 HCNN⫹O7HCN⫹NO 2.00 ⫻ 106 0.00 101 HCNN⫹O27O⫹CHO⫹N2 1.20 ⫻ 107 0.00 102 HCNN⫹OH7H⫹CHO⫹N2 1.20 ⫻ 107 0.00 103 HCNN⫹H7CH2⫹N2 1.00 ⫻ 108 0.00 104 N2O⫹HCN7HNCO⫹N2 1.20 ⫻ 107 0.00 105 HNCO⫹NO7NNH⫹CO2 2.00 ⫻ 107 0.00 106 HNCO⫹NH27NCO⫹NH3 3.70 ⫻ 106 0.00 2.00 ⫻ 107 107 HNCO⫹H7NH2⫹CO 0.00 108 OH⫹HCN7HNCO⫹H 1.98 ⫻ 10⫺9 4.00 8.59 ⫻ 106 0.00 109 NCO⫹H27HNCO⫹H 110 HNCO⫹M7NH⫹CO⫹M 1.10 ⫻ 1010 0.00 Enhanced third-body efficiencies: N2 (1.5), O2 (1.5), H2O (18.6), other (1.0) 111 HNCO⫹O7NCO⫹OH 2.20 2.11 112 HNCO⫹O7NH⫹CO2 9.60 ⫻ 101 1.41 6.40 ⫻ 10⫺1 113 HNCO⫹OH7NCO⫹H2O 2.00 3.10 ⫻ 1011 114 CH2⫹NO7H⫹HNCO ⫺1.38 115 HNCO⫹O7HNO⫹CO 1.50 ⫻ 102 1.57 116 HNCO⫹OH7NH2⫹CO2 3.30 1.50 3.80 ⫻ 107 117 CH2⫹NO7HCNO⫹H ⫺0.36 118 HCNO⫹H7HCN⫹OH 2.70 ⫻ 105 0.18 1.70 ⫻ 108 119 HCNO⫹H7NH2⫹CO ⫺0.75 120 HOCN⫹H7HNCO⫹H 2.00 ⫻ 101 2.00 121 HCNO⫹H7HNCO⫹H 2.10 ⫻ 109 ⫺0.96 122 HCN⫹OH7HOCN⫹H 2.00 ⫻ 101 2.00
0.00 0.00 0.00 0.00 0.00 0.00 10.00 12.56 4.19 37.68 360.00
[40] [40] [40] [40] thiswork thiswork thiswork [10] [10] [10] [42]
47.90 35.70 10.70 5.32 184.23 15.07 2.43 8.88 12.10 8.37 11.93 8.37
[42] [42] [42] [43] [43] [43] [43] [43] [43] [43] [43] [43]
of the analyzed mechanisms are unable to reproduce the measured [NO], because the calculated [NO] are approximately three times lower than the experimental ones, and only the fourth solid line for the modeling using the extended thermal mechanism plus reaction (10) describes well the experiments’ results. The results of the modeling using proposed mechanism are shown as solid lines in Figs. 2– 4. Thus, in all the flames analyzed, this mechanism
Fig. 1. Comparison of [NO] and [HCN] generated by the Fenimore mechanism (- - -) and as modeled by proposed mechanism via HCN (——).
has proved its ability to reproduce the experimental results for CH4 flames. The contribution of the key proposed reactions into forming NO via the inter-
Fig. 2. Comparison of the [NO] measured by Malte et al. [21] (Œ) within the premixed flames of CH4 ⫹ air and calculated by four mechanisms: (- - -) Miller and Bowman [10], (— 䡠 —) extended thermal mechanism (1-5) [11], (— —) extended thermal mechanism (1-5) plus reaction (26) from Bozzelli et al. [26] and (——) mechanism proposed here.
318
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Fig. 3. Comparison of [NO] measured by Sarofim and Pohl [22] (Œ) along a slightly rich premixed flame of CH4 ⫹ air (air ⫽ 0.95 ⫻ stechiometric) and modeled by four mechanisms: (- - -) Miller and Bowman [10], (— 䡠 —) extended thermal mechanism (1-5) [11], (— —) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (——)mechanism proposed here.
mediate HCN in rich flames of CH4 ⫹ air is higher than the contribution of Fenimore’s reactions. The extended thermal mechanism produces ⬃15% of the [NO] in this case. In lean CH4 ⫹ air flames, however, the most important contribution is from the extended thermal mechanism, in which the reaction (4) is decisive. In rich CH4 ⫹ air flames the proposed mechanism is the most important source of NO through HCN; in fact, the reactions initiated by N2O with
Fig. 4. Comparison of [NO] measured by Hayhurst and Hutchinsom [15] (Œ) along a rich premixed flame of CH4 ⫹ air (air ⫽ 0.67 ⫻ stoichiometric) and modeled by four mechanisms: (a) proposed mechanism, (b) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (c) Miller and Bowman [10], (d) extended thermal mechanism (1-5) [11].
Fig. 5. Comparison of [NO] measured by Hayhurst and Hutchinsom [15] (Œ) along a very rich premixed flame of H2 ⫹ air (air ⫽ 0.67 ⫻ stechiometric) and modeled by four mechanisms: (a) proposed mechanism, (b) extended thermal mechanism (1-5) plus reaction (6) from Bozzelli et al. [26], (c) Miller and Bowman [10], (d) extended thermal mechanism (1-5) [11].
CH3 in (65) are ⬃25 times more effective than the reactions initiated by NNH with CH3 in (66). The new mechanism generates much more NO than Fenimore‘s mechanism in the early stages of combustion, as seen in Fig. 1, because of the high [CH3] and [CH3O]. In this region, the reverse of reaction (67) is an additional source of N2O, without which the production of HCN via reactions initiated by N2O reacting with CH3 in (65) would die out. However, even in those parts of a flame where CH3O is almost fully consumed, the contribution of the proposed mechanism is higher than of Fenimore’s reactions because of more HCN being produced earlier.
Fig. 6. Comparison of [N2O] and [N2] measured by Vandooren et al. [23] along a CH4 ⫹ N2O ⫹ Ar flame and modeled by the proposed mechanism.
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Fig. 7. Comparison of [CH4], [CH3], and [NO] measured by Vandooren et al. [23] along CH4 ⫹ N2O ⫹ Ar flame and modeled by the mechanism proposed.
A crucial test of the proposed mechanism was the modeling of the experiment by Vandoren et al. [23]. This is presented in Figs. 6 – 8. Only in the case of NCO and N2 can some deficiency of the modeling be observed. It must be stated here that axial diffusion played an important role in this low pressure flame, without which mainly the position of the calculated curves would be delayed in comparison with the experiments. As can be seen HCN is formed in this flame according to the proposed mechanism and not by Fenimore’s reactions. The values of the rate coefficients for reactions (65) and (67) are crucial for modeling CH3. Without these reactions the maximum in [CH3] would be several times higher than measured by Vandooren et al. [23]. Figure 8 shows that the rate constants from Table I reproduced the measured maximum in [CH3].
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Fig. 8. Comparison of the [HCN], [NCO], and [HNCO] measured by Vandooren et al. [23] along the flame of CH4 ⫹ N2O ⫹ Ar and modeled by the mechanism proposed.
Simplified mechanism forming NO in flames of CH4 ⴙ air The detailed combustion mechanism [10,27] permits a full analysis of a flames structure and also a prediction of the concentrations both the radicals and the stable species in the flame. However, longer times are require for computing. For this reason the application is limited to only simple one-dimensional laminar flames or perfectly stirred reactors. For more a complex geometry in e.g. industrial diffusion flames, the calculations become very expensive. This economical problem may be solved using simplified mechanisms for both combustion and the formation of pollutants. Assuming that [NO] is far from equilibrium, as is often the case in practical applications, the following simplified formula may be used for the rate of forming NO:
d关NO兴/dt ⫽ 2k1f关N2兴关O兴 ⫹ 2k4f关N2O兴关O兴 ⫹ 2k6b关N2O兴关H兴⫹2k10f 关NNH兴关O兴 ⫹ 2k56f关N2兴关CH兴 ⫹ ⫹ 2k65f 关N2O兴关CH3兴, where kif and kib, respectively, denote the rate constants of the forwards and backwards steps of reaction i. The terms in Eq. (I) represent the rate of NO formation from: Zeldovich’s reactions (1) and (2) with the assumption of a steady state for atomic nitrogen [1]; from N2O via (4) and the reverse of (6); the fourth term represents NO from NNH; the fifth term is Fenimore’s mechanism and finally are the reactions initiated by N2O reacting with CH3, i.e. the new mechanism. It was assumed that NH,
(I) N, and HCN are all fully converted to NO [28,29]. In and near the reaction zone, where large concentration of the free radicals H and OH are present, the N2O formed in the reverse of reactions (3) and (67) may be destructed not only by reactions (4), (5) and (65), but also by reactions (6), (7) and (8). Thus [N2O] within a flame quickly reaches a semisteady state at ⬃1 ppm. In accordance with the steady state assumption, [N2O] can be calculated from:
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关N2O兴 ⫽ 兵k3b关N2兴关O兴关M兴 ⫹ k67b关CH3O兴关N2兴其/兵k3f关M兴 ⫹ 共k4f ⫹ k5f兲关O兴 ⫹ (k6b ⫹ k7f兲关H兴 ⫹ ⫹ k8f关OH兴 ⫹ 共k65f ⫹ k67f兲关CH3兴} Glarborg et al. [30], on the basis of experiments conducted at temperatures below 1400 K, suggested that reaction (8) may play a minor role destroying N2O. They found that its rate constant is ⬃10 times lower than the values recommended by Miller and Bowman [10] and ⬃60 times lower than the values recommended by Tsang and Herron [31]. A similar conclusion has been reported by Sausa et al. [32]. Despite this, reaction (8) was included in Eq. (II). In accordance with Bozzelli and Dean’s [13] suggestion, [NNH] can be calculated assuming that the fast reaction (9) is equilibrated, so [NNH] ⫽ [N2] [H]/K9, where K9 is the equilibrium constant of (9). Application of the above equations requires a knowledge of [O], [H], [OH], [CH], [CH3], and [CH3O]. Equation (I) is able to predict well the rates of formation of NO in a flame, if the concentrations of the main radicals are properly established. To simplify calculating the concentrations of the radicals, many ways based on partial equilibrium and the steady state assumption were proposed [22,29,33– 36]. The difference between the detailed mechanism in Table I and the above simplified formulae, which require the rate constants for only 12 reactions: (1), (3–10), (56), (65) and (67), depends on the accuracy of these simplified calculations, of which the least accurate are for [CH3] and [CH3O]. The presence of methane within a flame creates problems for calculating [O], [H] and [OH] using partial equilibrium in both the systems of H2/O2 and CO/H2/O2 as already addressed [29] and demonstrated in Figs. 9 and 10 for of distribution of [O] and [OH] in a premixed, moderately lean (air ⫽ 1.1 x stoichiometric amount) flame of CH4 ⫹ air at constant temperature of 1800 K. The results obtained from the classical partial equilibrium assumption of H2/O2 [22,33] agree very well with the ones calculated using the detailed combustion mechanism, but only in the region where CH4 is completely consumed. Calculations at lower temperatures of 1700 and 1600 K gave similar results. Some of the proposed reduced mechanisms of combustion allow one to compute [H] [34,35,37]. In this case, if the local [H] in the flame is known, [O] and [OH] can only be obtained from assuming partial equilibrium of H2/O2 [35]. In contrast with the classical H2/O2 partial equilibrium, the [O] and [OH] obtained by this approach agree very well with those calculated from the detailed combustion mechanism. Figure 11 compares the rate of formation of NO and also [NO] modeled in the same CH4 ⫹ air flame (air ⫽ 1.1 x stoichiometric, T ⫽ 1800 K) as derived
(II)
Fig. 9. Modeled [O] and [CH4] in a flame of CH4 ⫹ air (air ⫽ 1.1 ⫻ stoichiometric, T ⫽ 1800 K): (——) [O] from detailed combustion mechanism, (— —) [O] from classical H2/O2 partial equilibrium, (- - -) [O] from H2/O2 partial equilibrium with [H] obtained from detailed combustion mechanism.
from the detailed computations for both combustion [10] and the production of NO, with the mechanism in Table I and secondly as derived from Eq. (I) using the mentioned two above-ways of assuming partial equilibrium of H2/O2. For hydrocarbon radicals in Eq. (I), expressions based on the steady state assumption were used: these were a slightly modified Eq. (3) in [36] for [CH3] and Eq. (44) in [29] for [CH]. The
Fig. 10. Modeled [OH] and [CH4] in the same flame as Fig. 9: (——) [OH] from detailed combustion mechanism, (— —) [OH] from classical H2/O2 partial equilibrium, (- - -) [OH] from H2/O2 partial equilibrium with [H] obtained from detailed combustion mechanism.
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assumed to be fast. Further investigations of the proposed reactions are necessary in order to improve the modeling of NO in hydrocarbon flames. The simplified mechanism (Eq.I) for the formation of NO, together with the assumption of partial equilibrium for [NNH], [O] and [OH] and steady state for [N2O], [CH], [CH3] and [CH3O] was found to be a good approximation for modeling [NO] in industrial flames of CH4 ⫹ air.
References Fig. 11. NO formation rate and [NO] in the same flame as Fig. 9 calculated from the detailed combustion and NO mechanisms (——) and also the simplified NO mechanism (Eqs.I-II) with [CH], [CH3] and [CH3O] from the steady state assumption and: (— —) [H], [O], [OH] from the classical H2/O2 partial equilibrium, (- - -) [O], [OH] from H2/O2 partial equilibrium and [H] from the detailed combustion mechanism.
steady state assumption was used also for calculating [CH3O]. As is seen from Fig. 11 the rates of formation of NO obtained with the classical H2/O2 partial equilibrium are very high, probably as a consequence of [O] being too large. However, in the case when only [O] and [OH] were evaluated using partial equilibrium of H2/O2 with [H] from the detailed combustion mechanism, the rates of formation of NO obtained from Eq. (I) are very close to those calculated from the detailed mechanism for combustion and NO. At short times, the simplified mechanism overpredicts [NO] due to the fact that formation and conversion of HCN into NO are assumed to proceed simultaneously, while in the detailed mechanism a considerable delay time can be noticed. However, the final [NO] is smaller because of this simplification.
Conclusions The proposed reactions of N2O and NNH with CH3 radicals provide an additional mechanism forming NO in hydrocarbon flames. The role of this mechanism depends on the rate constants of the proposed reactions, which at present could only be estimated. Nevertheless, the rate constants of the new reactions proposed in Table I result from an analysis of experimental measurements. Introduction of new reactions always creates problems, but all the experiments analyzed [15,21– 23] could only be described if the new reaction N2O ⫹ CH3 3 CH2NH ⫹ NO and the previously studied reaction N2O ⫹ CH3 7 CH3O ⫹ N2 [24,25] were
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