Kinetics and mechanism of NH3 oxidation

Nineteenth Symposium (International) on Combustion/The Combustion Institute, 1982/pp. 97-105

KINETICS

AND

MECHANISM

OF NH 3 OXIDATION

A. M. DEAN, J. E. HARDY, XND R. K. LYON Corporate Research--Science Laboratories Exxon Research and Engineering Company Linden, New Jersey 07036 The kinetics of NH3 oxidation at 1279 to 1323 K have been studied experimentally with a flow tube reactor and theoretically by computer modelling. It was observed that the rate of NHa oxidation was zero order in NH3, increased with increasing [O2], was inhibited by the presence of initial water, and promoted by the addition of H2 or NO. During NH3 oxidation, [NO] increased rapidly initially and more slowly thereafter. The size of the initial rise was independent of the presence or absence of H20, while the subsequent increase was reduced by H20. These and other kinetic features including the rate of oxidation of a trace of CO (a monitor of [OH]) are all quantitatively explained by a computer model. The major features of this model are conversion of NH 3 to NH2 via reaction with O and OH, the reaction of NHz with 02, O, and OH to ultimately form NO, and the concurrent reduction of NO by NH2 to ultimately yield N 2. The chain carrier concentration during the process is regulated by a balance between carrier formation via the chain branching reactions and their removal by reactions such as NH 2 + HNO ~ NH3 + NO.

these simplifying conditions, it was possible to describe quantitatively the observed NO reduction kinetics via a simple, plausible mechanism. Su bsequently, Silver, Gozewski and Kolb, (6) Miller, Branch and Kee,/7) and Salimian and HansonIs) attempted the more difficult task of modelling the kinetics of the N H 3 - - N O - - O 2 system from 1000 to 1500~ K, i.e., modelling for conditions under which both NH3 reduction of NO and oxidation to form it are important. This required adding a great many reactions to the simple mechanism of Reference 5, making it so complex as to be tractable only by computer. Although the schemes produced by the three groups of workers differed in many details, they are successful in their stated purpose: providing a mechanism that qualitatively modeled the temperature behavior of the Thermal DeNO x chemistry. If, however, one attempts to predict the kinetics of NH 3 oxidation via these mechanisms, the resultant predictions are significantly different from the observations reported below. For example, the mechanism of Miller et al., while correctly predicting that the NH 3 decay profile is linear, also predicts much too a fast a rate of NO production and completely fails to predict the change in the character of the NO production which we observe. The mechanisms of Salimian and Hanson and of Silver et al. are generally quite similar and specifically both fail to incorporate adequate chain terminating reactions. Calculations with the latter model show ammonia oxidation under the

Introduction

Although the kinetics and mechanism of the homogeneous gas phase oxidation of NH 3 have been studied by a number of investigators,/1-31 the mechanism of the reaction is still in considerable doubt. This is unfortunate both because of the reaction's scientific interest and because of its practical importance. For example the onset of NH 3 oxidation imposes an upper limit on the tempera. (4) ture range for whieh the Thermal DeNOx process may be used to reduee NOx emissions. As has been diseussed previously,/5/NH3 oxidation involves two sets of reactions occurring eoneurrently. In one set, NH 3 oxidizes to form NO while, in the other, it reduces NO to N2. Consequently, during NH 3 oxidation, NO tends toward something resembling a steady-state concentration. In a narrow range of temperature centered at 1250 K, the overall reaction is rapid and the pseudo-steady-state concentration of NO is low. Thus one can, as is done in the Thermal DeNO~ process, inject NH3 into the NO-eontaining flue gas in a boiler or furnace at a point where the flue gas temperature is near 1250 K and reduce most of the NO to N2 via a homogeneous gas phase reaction. In a previous study,/5/ one of us examined the kinetics of the NH3-NO-O~ reaction system at temperatures centered around 1170 K, i.e., conditions sueh that NH 3 reacts primarily by reducing NO, oxidation to form NO being a minor process. Under 97

98

REACTION MECHANISMS AND MODELING I

conditions applied here to be an explosive process, in contradiction to our experiments. The present study deals, experimentally and by computer modelling, with the kinetics of NH 3 oxidation in the temperature range 1279 to 1323 K. It is the first study to provide a satisfactory, detailed, quantitative account of NH z oxidation.

Experimental Method Since our flow reactor and experimental procedures have been previously described(5'9) only a brief description will be given here. Flowing mixtures of the reactant gases in He carrier gas were prepared with specifically calibrated rotameters. Water was metered via a Cheminert metering pump and was added to the flowing gas via an evaporator packed with 2 mm glass beads. The gaseous mixture then went to a tubular quartz flow reactor: in most experiments a reactor of conventional design---0.2 cm i.d. • 45 em inlet/ heatup zone, a 1.0 cm • 45 cm reaction zone and a 0.2 cm i.d. • 45 cm outlet/quench zone, but in other experiments a "folded" reactor with a 0.2 cm i.d. • 40 cm inlet followed by a folded section of 5 pieces of 0.2 cm i.d. • 50 cm each and coneluded by a 0.2 cm i.d. • 40 cm outlet. The latter design has the advantage that the time for diffusion from a point equidistant from the eenterline and wall was only 0.00013 seconds (estimating DNH3 He at 1300~ K as 9.6 cm2/sec via Brokaw's(~4) technique). Since this diffusion time is short compared to the reaction time, plug flow is assured. Similarly, for the conventional reactor, the estimated diffusion time becomes 0.0032 seconds,, again much shorter than the reaction time, assuring plug flow. This analysis is confirmed by the observed agreement of NH 3 decay rates in the two reactors. These reactors were housed in a three zone electric furnace whose temperature was monitored by a set of chromel-alumel thermocouples, these thermocouples being placed so that they sensed the reactor wall temperature. Heat transfer calculations assured us that the helium used as carrier gas in these experiments would come to within 1~ of the reactor wall temperature during its passage through the inlet zone. The reactor temperature was controlled to be constant and the same within all three zones of the electric furnace to within 1~ C. After passage through the reactor, the gases went to analysis, HzO being measured by a DuPont 510 moisture analyzer and CO z by a Beckman Model 865 infrared analyzer. The latter instrument was also used to measure CO via the expedient of passing the gas mixture through Mallcosorb to remove CO 2 and then converting the CO to CO 2 by catalytic oxidation. NO and NH3 were measured with a Thermoelectron 10A chemiluminescent analyzer via the technique of Hardy and Knarr. (1~

Results Experiments were done, observing both NH 3 decay and NO formation as functions of reaction time for XNH3~ = 900 ppm, Xo~~ = 2, 4 and 8%, Xnzo, = 0 and 1%, for temperatures ranging from 1279 to 1323 K at a total pressure of 120 kPa for a variety of surface to volume ratios, and for various pretreatments of the surface of the quartz reaction vessel. In these measurements extreme care was necessary to obtain meaningful absolute rate data. In addition to these absolute rate measurements, considerable data were collected with sufficient care to provide reliable relative measurements of the effect of various factors upon the reaction rate. All experiments were done with helium carrier gas. It was found that, after an induction period, NH 3 decayed linearly with time, i.e., the decay was zero order in NH 3. During NH z oxidation, NO is produced rapidly at the start of the reaction and more slowly thereafter. The rate of NO production after the initial burst is significantly greater for initially dry NH 3 than for wet NH3, as is illustrated by the data i n Figures 1 and 2. Similarly, addition of

d-

o d u o , , oo

Z d(a) !

~NO

i

!

i

!

!

~

(bJ

o

#

0

s'o TIME (msec)

FIG. 1. Concentration-time profiles for T = 1279 K, 4% Oz, and no added H20. The dashed lines are least square fits to the observed data, and the solid lines are calculated using the mechanism shown in Table III.

KINETICS AND MECHANISM OF NH3 OXIDATION

Computer Modelling: Procedures and Results

"'--~.. ""~....

O-

Zo-

-1z r 6. (a)

o

E

o

/

Io

o

99

(b) I

i

!

f

i

so

1oo

Iso

2oo

2so

3oo

TI ME (msec) FIG. 2. Concentration-time profiles for T = 1279 K, 4% O2 and 1% H20. Dashed lines are least square fits to the data, and the solid lines are calculated from Table III. water inhibits the ammonia decay rate. Varying the reactor surface to volume ratio from 3.99 cm -1 to 20 cm -1 and changing the quartz surface from one that had seen extensive service, to one freshly fabricated, to one washed with chromicsulfuric acid cleaning solution caused significant variations in the induction time but had no influence on either the post induction time NH 3 decay rate or on the NO production rate. It was found that the post induction time kinetics could be accurately characterized by three parameters: (1) --d(XNHJXNn3,)/dt, the (constant) fractional ammonia decay rate, (2) dXno/dt, the (constant) rate of NO increase at later times and (3) NOI, the value of XNO obtained by extrapolation to t = 0 of the later time NO profile, i.e. NOi characterizes the initial burst of NO production. These kinetic parameters are listed in Table I. Experiments also were done under conditions given in Table II in which NH 3 was oxidized in the presence and absence of trace CO. The rate of NH 3 oxidation was unaffected by the CO, while the rate of concurrent CO oxidation was as shown in Table II. This table also shows the results of experiments in which the NH 3 oxidation was perturbed by adding small amounts of NO or H 2.

The fact that the three kinetic parameters were all observed to be surface independent, while the induction time was surface dependent, implied to us that the reaction was, in part, surface initiated but that, after initiation, a branching chain mechanism became dominant, the kinetics of this mechanism being surface independent. Accordingly, in developing a mechanism which would account for our results, we attempted to model -dXNm/dt, dXNo/dt and NOi quantitatively while treating t i as merely an upper limit. Table III lists various elementary reactions, some accurately characterized and others speculatively proposed in the literature, all of which might be of importance to the oxidation of NH3. Such a mere listing of reactions, of course, says little about the reaction mechanism. To propose a reasonable mechanism, one must propose a set of rate constants for the reactions which is a priori plausible, consistent with previous knowledge, and which, when fed into a chemical kinetics computer program, produces predicted results consistent with observation. Given the complexity of the set of reactions in Table III, it may appear to the reader that adjusting the rate constants to fit the data is a difficult mathematical undertaking but one certain of ultimate success because so complex a reaction scheme must be capable of fitting any data set. The opposite, however, is the case. The reaction scheme in Table III is actually quite inflexible with few adjustable parameters. Thus, the task of trial and error fitting of the rate constants to the data is mathematically straightforward but that fit is not likely to be good unless the reaction mechanism is a reasonable approximation to reality. One reason for this situation is that the 31 reactions in Table III, 18 (Nos. 1, 2, 3, 4, 16, 18, 19, 20, 21, 22, 23, 24, 25, 26, 27, 28, 29, and 30) all appear to have accurately measured rate constants and thus cannot be used as adjustable parameters. Table III also shows the results of sensitivity calculations, calculations of the effect' on the observable kinetic parameters of varying the rate constants of the elementary reactions. Only four (#3, 4, 16 and 19) of the eighteen measured reactions turn out to be sensitive, i.e., if the literature value of these rate constants were in error, this would significantly influence the calculated results. Furthermore, these four reactions are quite insensitive relative to the asterisked reactions in Table III, and thus one confidently uses the literature values for all eighteen of these reactions. Similarly, in view of the calculation's lack of sensitivity to reactions 5, 14 and 17, these may be accepted at their literature values. Thus it appears that one has ten reactions (Nos. 6, 7, 8, 9, 10, 11, 12, 13, 15 and 31) with poten-

100

REACTION

MECHANISMS

AND

MODELING

I

TABLE I Summary of the measured kinetic parameters X~na, = 9 0 0 p p m P = 1 2 0 k P a

--d(XNHJX~Ha,)

dNO - -

(s -t)

dt T(K)

Xo2,(%)*

XH~O, = 0

1279

2.0 2.0

1 . 5 4 -+ . 0 5 1.64 • .10 --2 . 3 1 --- . 1 7 2.38 • .10 2 . 5 2 -+ . 0 4 2 , 4 3 -+ . 0 4 --4 . 0 0 -+ . 1 3 --3 . 3 7 -+ .21 6 . 8 3 -+ . 7 0 1 0 , 3 -- 3 . 3 6 . 4 5 -+ . 7 9

2.0

1312

1323

2.0 4.0 4.0 4.0 4,0 4.0 8.0 8.0 8.0 8.0 2.0 4.0 8.0 2.0

(ppm s-')

NO, (ppm)

dt Xuzo, = 1 . 0 % --0.83 0.99 • ---1.88 • 1 , 7 0 -+ 3.29 • -2.76 • -3.15 • 5.65 • 9 . 4 4 -+ 5.00 •

.08 .02

.07 .15 .13 .13 .21 .60 3.3 .38

Xa2o, = 0

XHZO,= 1 . 0 %

--0

--

~0

--

--~7 --2 0 , 0 -+ --6 9 -+ --5 4 -+ 151 -2 5 1 -2 0 8 -+

2

6 8 75 12

Xazo, = 1 . 0 %

5 10

~0 ~15 -2 0 --- 3 -5 . 6 -+ 1 . 4 ---19 -+ 5 4 8 -+ 6 2 8 -+ 4 3 8 -+ 9 8 7 -+ 31 6 8 -+ 5

,2

Xn~o, = O

--

4

--

2

8 7 . 0 -+

--

7 . 1 -+ .3

.5

6 , 9 -+ , 3

2 0 . 6 --- .2 --

1 3 -+

1

--

14-+1 5 . 5 -13 • 2 8 -+ 7 . 6 -+

.8 1 4 .9

4 8.2 28 1.5

-+ -+ -+ -+

1 1.2 8 1.6

* M u l t i p l e e n t r i e s f o r t h e s a m e Xo2~ a r e r e s u l t s w i t h d i f f e r e n t r e a c t o r s ( s e e text).

T A B L E II N H a o x i d a t i o n in t h e p r e s e n c e

Observed

A d d e d gas 1. C O / H 2 0

2. N O

relative rates

d ( l n C O / C O , / dt)H~O. . . . o~

P = 120 kPa XNH3, = 9 0 0 p p m Xco~ = 18 p p m Xoz, = 4 %

d ( l n C O / C O , / dt)n2o, = o z~

T = 1312 K

- ( d N H f f d t ) N o , = 9s ppm

T = 1312 K P = 120 k P a XNn3, = 9 0 0 p p m Xo2, = 2 %

*Using mechanism

gases

T = 1270 K

P = 120 kPa XNH3, = 9 0 0 p p m Xo2, = 2 % 3. Hz

of added

Calculated relative rates* 1.75

= 1.8

-(dNHffdt).~o, = o

1.51 = 1.78

- ( d N H f f d t ) n 2 , = ~spp.1 -(dNH3/dt) .... 0

a n d r a t e c o n s t a n t s l i s t e d in T a b l e III.

1.2 1.3

KINETICS AND MECHANISM OF NH3 OXIDATION

101

TABLE III Proposed mechanism for NH3 oxidation at T -~ 1300 K and sensitivity results k = AT" e -E/RT (Units are mole, cm 3, sec)

Reaction 1. 2. 3. 4. 5, *6. *7. *8. *9. *10. "11. '12. '13. 14. *15'. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. '31.

NH~ + M = NH2 + H + M NH3 + H = NH2 + Hz NH3 + O = NH2 + OH NHa + OH = NH~ + HaO NH2 + H = NH + H~ NHa + O = NH + OH NHz + OH = NH + HaO NH2 + 02 = HNO + OH NH2 + NO = NNH + OH NH2 + NO = N~ + H20 NHa + HNO = NH3 + NO NHa + NNH = NH3 + N~ 1%NH + M = Na + H + M NNH + OH = N2 + H20 HNO + M = H + NO + M HNO + OH = NO + H20 NH + Oa = HNO + O Ha + OH = H + H20 H + Oz = OH + O O + Ha = OH + H OH + OH = O + HaO H + O z + M = HO~ + M H + HO~ = OH + OH O + HO2 = O~ + OH OH + HOa = HaO + O2 NO + HO 2 = NO2 + OH H + NO~ = NO + OH O + NO~ = NO + O2 NO2 + M = NO + O + M O + O + M = O2 + M NNH + NO = N2 + HNO

log A

n

16.68 13.40 12.18 12.52 11.70 13.30 10.48 13.71 19.79 19.96 14.26 13.00 16.86 13.48 16.28 13.56 13.48 13.34 14.34 10.26 12.80 15.18 14.40 13.68 13.70 12.53 14.54 13.00 16.04 18.15 12.40

0 0 0 0 0.50 0 0.68 0 -2.46 -2.46 0 0 0 0 0 0 0 0 0 1.00 0 0 0 0 0 0 0 0 0 -1.00 0

Sensitivity of Kinetic Parameters

Conditions of Figure 1 -d(NH3/ ReferE / R • 10 -3 NH3~)/dt NO~ dNO/dt ence 47.3 8.6 3.0 1.1 1.0 0.5 0.6 15.1 0.9 0.9 0.5 0 22.6 0 24.5 0 1.7 2.6 8.5 4.5 0.5 -0.5 1.0 0.5 0.5 -0.1 0.8 0.3 33.2 0.2 0

0.0 0.0 -0.032 0.021 0.025 0.10 0.072 0.52 0.45 -0.44 -0,57 -0.18 0.19 -0.036 0.50 -0.036 0.013 0.0 0.13 0.0 -0.01 -0.032 0.0 0.0 -0.027 -0.022 0.0 0.0 0.0 0.0 -0.041

0.0 0.0 -0.046 -0.046 0,0 0.045 0.0 0.46 -0.28 -0.16 0.14 0.0 -0.046 0.0 -0.081 0.045 0.0 0.0 -0.037 0.0 0,0 0.0 0,0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -0.018

0.0 0.0 -0.22 -0.12 0.0 0.50 0.27 1.41 0.77 -1.50 -1.19 -0.40 0.45 0.0 1.12 -0.17 0.0 0.0 0.32 0.0 -0,056 -0.051 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -0.046

11 12 13 6 18 19,20 21 Adj 6 6 Adj 7 Adj 7 13 14 7 14 14 14 14 14 14 15 15 16 13 13 13 17 Adj

*Kinetically sensitive reactions, see text. NOTE: Each sensitivity is the log of the ratio of the kinetic parameter with k = 2 • standard to this same parameter with k = 0.5 • standard. A value of 0 means that this factor of four change i n k had no effect. Values less than zero imply inhibition, while those greater than zero imply promotion.

tially adjustable rate constants, but further examination reveals additional constraints. Reactions 6 and 7 are expected to be rapid./19'2~ The k6 value in Table III is consistent with the room temperature measurements of Albers, et al. (19/ and Gehring et al. (20/ and k7 was taken from the literature. (21) Reactions 9 and 10 have been measured in the temperature range of interest by Silver, et alJ 6) and the fitted values in Table III agree with the observed ratio kg/klo with the rate constants 30% higher than reported. A value for kt2 was

taken from the literature, IT> a n d the value used from k15 is consistent with lower temperature measurements of k_ls. (13) Thus only the rate constants of reactions 8, 11, 13 and 31 were truly adjustable parameters and the authors believe that the fitted values in Table III are all within theoretically plausible bounds. Typical fits of the mechanism to our data are shown in Figures 1, 2 and 3 and Table II. Not surprisingly, in view of the heterogeneous effects discussed above, the experiments show shorter indue-

102

REACTION MECHANISMS AND MODELING I

m

"I-

z~. i"0

I m §

~" 20v

~.

15~

P

0 z

5,

60"

Q.

30 z -o

63

15. V i (

PERCENTO X Y G E N

FIc. 3. Kinetic parameters for the dry case at 1279 K. V, A, +, [] = different reactors. 9 = calculated using Table III. tion times than do the calculations. The agreement between calculation and observation with respect to NH 3 decay and NO formation after the induction time is excellent, especially so since this agreement persists when the system is perturbed by varying the oxygen concentration, going from dry to wet, and adding NO and H z. Further, the calculation accurately predicts the rate of CO oxidation when small amounts of CO are cooxidized with NH 3. This is quite important, since CO is oxidized almost exclusively by OH and thus accurate prediction of the CO oxidation rate implies that the calculated [OH] is also accurate.

NH a is converted to NHz, which during the early stages of reaction reacts with 02 to form HNO ultimately producing NO, and which at later times also reacts with OH and O, again ultimately producing NO. Along the other path NH3 forms NH2 which reacts with NO ultimately yielding Nz. Both channels involve both chain branching and chain terminating steps with the result that instead of growing explosively the chain carrier concentration stabilizes at quite finite values. Some further insight may be gained by oversimplifying the mechanism to the point that it becomes analytically tractable. Let us neglect all reactions except reactions 3, 4, 8, 9, 10, 11, 13, 15 and 19 and make the usual steady state assumptions. At steady state the total rate of chain carrier production via the chain branching reactions is equal to the rate of chain carrier removal via the termination reactions. It is convenient to do the accounting for the various chain carriers in terms of NH 2, i.e., an OH which reacts to yield an NH 2 is worth one NH2, an O which reacts to yield an NH2 and an OH which yields another is worth two NH2s, etc. It is also convenient to define ct as ct = kg/(k9 + k10). By examining the reaction set one can readily satisfy oneself that Reaction 15 is the critical chain branching step since without it the concentration of all chain carriers except HNO tends toward zero and HNO is no longer a true chain carrier but rather a stable product. Thus one may write the balance between chain carrier production and removal as ~Ikls[HNO][M] = 13kll[NH2][HNO ], where 13 and ~ are stoichiometric coefficients, the number of NH2s or their equivalent which are produced or destroyed by the occurrence of reactions 15 and 11 respectively. Reaction 15 yields an H and an NO, the H reacting to produce an OH and an O, and thus being worth 3 NH2s and the NO consuming an NH2 to yield arts and otOHs. Conse= quently ~ = 3 - 1 + 3~ + ot = 2 + 4et. Reaction 11 consumes one NHz directly and by producing an NO consumes 1 - 4a NH2s indirectly, which means than 13 = 2 - 4or. If one uses the steady state equations implied by d[NHz]/dt = 0, d[HNO]/ dt = 0, and d[NO]/dt = 0, it can readily be shown that d[NHa]/dt = -2ks[NH2]ss[O2], which means that d[NH3]/dt = - ( 2 + 4a)kskls[Oz][M]/ (1 - 2ct)kll. N2+H20

N ~ '

/

Discussion

While Table II1 formally defines a reaction mechanism whose kinetic predictions match experiment, the reader may have trouble "seeing the forest from, the trees." Figure 4 shows an overview of the principal reaction pathways. Along one path

NH3 O~ NH2

N~

NH3+N2

-OH+NNH ~

N2 CHANNEL

N2+H

OH~'~ ~ NH3 + ND NH~22HNO ~-~ H+NO ~H H20+NO

NO CHANNEL

FIG. 4. Overview of NHa oxidation near 1300 K.

KINETICS AND MECHANISM OF NH3 OXIDATION The simple mechanism thus predicts that NH 3 oxidation will be zero order in NH 3 as was observed and first order in 02 , in rough accord with experiment. It can also readily be shown that [NO]ss = ks[O2]/(k 9 + klo) i.e., the simple model predicts that during NH3 oxidation NO will rapidly rise to an initial value, as is observed, and hold constant thereafter, in disagreement with observation. If one wishes to correct this defect in the simple model, one must allow NH2 to oxidize not only by reaction with 02 but also by reaction with 0 and OH (reactions 6 and 7 followed by 17). While this makes the model too complex to be analytically tractable, the computer model shows the effect of this change. [NH2] is no longer constant; instead the balance between chain carrier formation and destruction maintains a constant rate of NH 3 oxidation by controlling the overall chain carrier concentration while allowing the proportion of NH 2 to OH and O to decrease as the reaction proceeds, which causes NO to increase. While the fact that H20 addition, which increases [OH], retards NH 3 oxidation may at first glance be surprising, it is readily understandable in view of the fact that O is more effective for oxidizing NH 2 than is OH and that H20 addition suppresses [O] via the reverse of Reaction 21. Because the mechanism in Table III provides an accurate and readily understandable account of the complex kinetic behavior of NHa oxidation, the authors believe that it must, fairly closely, resemble the actual reaction mechanism. It is, however, important to recognize a number of ambiguities. The numerical values for the four adjustable rate constants given in Table III are the result of trial and error fitting and are not in any sense a unique set. Further, while some of the reactions in Table III appear unambiguously necessary if one is to construct a mechanism that fits the data, others could be modified without altering the fit. For example, the products of Reaction 6 are ambiguous. As written, it is NH 2 + O ~ NH + OH, the sole fate of the NH being to react with 02 forming HNO and O. It could, however, equally well be NH2 + O ---) HNO + H, since the predominant fate of H is Reaction 19, H + O2 --~ OH +O, The dominant fate of NzH is Reaction 13, hence with some readjustment of rate constants Reaction 9 could be written as NH 2 + NO ---) N2 + H + OH instead of NH 2 + NO ---) N2H + OH. The results of Silver(22) and Wolfrum(z3y, however, both strongly indicate that the NH 2 + NO reaction does not directly produce H atoms. Finally, Reaction 8, NH 2 + O 2 ---> HNO + OH, could also be written as Reaction 8', NH 2 + O 2 --> HNOOH, the HNOOH being consumed by competing Reactions 15', HNOOH ~ H + NO

103

+ OH, and 16', HNOOH + OH ---) HzO + NO + OH, and 11', HNOOH + NH z---> NHa + NO + OH. Provided that ks, = ks, kls, = k15 [M], kls, = kls, and ku, = ku, these different reaction pathways would give identical fits to our observations. One could, of course, distinguish these pathways by doing experiments in which [M] was varied. While detailed discussion is beyond the scope of this paper, the authors have attempted such experiments and the preliminary results seem to favor Reaction 8'. Despite this ambiguity in the products of Reaction 8, the authors do believe that the reaction itself is quite unambiguous, i.e., there seems no other plausible way to account for the early production of NO. Reactions involving OH and O cannot be responsible since water addition changes the concentrations of these species but does not change [NO] v Furthermore, [NO]/is most sensitive to ks, k9, and klo. To the extent that k9 and klo are known from the literature, the present study defines the value of ks .

Acknowledgment The authors wish to thank David Stern for his skilled assistance.

REFERENCES 1. (a) D. C. BULL, Combust. Flame 12, 603 (1968) and references therein. (b) N. Fujn, H. MIYAMA, M. KOSHI AND T. ASARA, Eighteenth Symposium (International) on Combustion, Combustion Institute, Pittsburgh, PA, 1981, p. 873. 2. W. E. KASKANAND D. E. HUGHES, Combust. Flame, 20, 381 (1973) and references therein. 3. P. B. PABSBERG, J. ERIKSEN AND H. C. CHRISTENSEN, J. Phys. Chem., 83, 582 (1979) and references therein. 4. (a) R. K. LYON, U.S. Patent No. 3,900,554; 1975. (b) R. K. LYON, Hydrocarbon Processing, 109-112 (October 1979). 5. R. K. LYONAND D. J. BENN, Seventeenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, PA (1979), p. 601. 6. J. A. SILVER, C. M. GOZEWSKIAND C. E. KOLB, ARI-RR-66, Western States Section/Combustion Institute, Fall 1980. 7. J. A. MILLER, M. C. BRANCHAND R. J. KEE, SAND 80-8635, Sandia Laboratories, 1980. 8. S. SALIMIANAND R. K. HANSON, Combust. Sci. Technol., 23, 225 (1980). 9. R. K. LYON, J. E. HARDY,AND D. J. BENN, Fall 1978 Meeting/Western States Section of the Combustion Institute.

104

REACTION MECHANISMS AND M O D E L I N G I

10. J. E. HARDY AND J. J. KNARR, Journal of the Air Pollution Control Association, in press. 1i. K. HOLZRICHTER AND H. GG. WAGNER, Eighteenth Symposium (International) on Combustion, C o m b u s t i o n Institute, P i t t s b u r g h , PA (1981), p. 769. 12. J. E. DOVE AND W. S. NIP, Can. J. Chem., 52, 1171 (1974). 13. D.. L. BAULCH, D. D. DRYSDALE AND D. G. HORNE, Evaluated Kinetic Data for High Temperature Reactions, Vol. 2, Butterworth, London (1973). 14. D. L. BAULCrl, D. D. DRYSDALE, D. G. HORNE AND A. C. LLOYD, Evaluated Kinetic Data for High Temperature Reactions, Vol. 1, Butterworth, London (1972). 15. A. C. LLOYD, Int. J. Chem. Kinetics, 6, 169 (1974). 16. M. T. LEU, J. Chem. Phys., 70, 1662 (1979). 17. H. S. JOHNSTON, NSRDA-NBS, 20 (1968).

18. G. S. BAHN, Reaction Rate Compilation for the N-O-H System, Gordon and Breach (1968). 19. E. A. ALBERS, K. HOYERMANN, H. GG. WAGNER AND J. WOLFRUM, Twelfth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, PA (1969), p. 313. 20. M. GEHRING, K. HOYERMANN, H. SCHACKE AND J. WOLFRUM, Fourteenth Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, PA (1973), p. 99. 21. R. TUNDER, S. MAYER, E. COOK AND L. SCHIELER, Aerospace Corporation Report No. TR1001 (9210-02)-1, AD 813 485 (1967). 22. J. A. SILVER, private communication. 23. J. WOLFRUM, private communication. 24. R. S. BI~OKAW,Alignment Charts for Transport

Properties--Viscosity, Thermal Conductivity, and Diffusion Coefficients for Non-Polar Gases and Gas Mixtures at Low Density, NASA TR R-81 (1961).

COMMENTS S. J. Harris, General Motors Research Labs, USA. How strongly correlated are the values you obtained for the 4 unknown rate constants? That is, if you arbitrarily change the value of one of them, can the data still be fit by adjusting the others?

Author's Reply. That depends on the reaction. The NH~ + 02 reaction, for example, very strongly influences t h e predicted value of NO~, the NO produced at the start of the reaction. Since none of the other reactions have nearly as much effect on NO,, it would be hard to fit this aspect of our data without the right rate constant for NH~ + O~. On the other hand the N~H + NO reaction is only marginally sensitive. We could change it, readjust the other rate constants and come out with a fit that's only slightly poorer. $

J. A. Miller, Sandia National Laboratories, Livermore, USA. I am curious about the possibility of forming NO2 in your experiments. Did you look for NO2 and did you detect any? There should exist the possibility of NO2 formation from the reaction sequence 0 2 - - * N z + HO2

(i)

HO2 + NO --) NO2 + OH.

(ii)

NNH§

If you did not detect NO2, what does that imply about an upper limit for the rate coefficient of reaction (i)? In addition, I would like to reply to Professor Kaufman's remark yesterday that your activation energy for the reaction N N H + M--~N 2 + H + M is too high. I also believe it is too high, but I think the A factor is also about two orders of magnitude too high, perhaps compensating for the activation energy at the temperature of your experiments. One can estimate a strong collision A factor for a weakly bound triatomic molecule to be about 1016 cm3/mole-sec, very low density of states at the dissociation limit (see Troe's 1977 J. Chem. Phys. papers). One can also infer from Troe's papers that 13c should be of the order of .01 or .02 and that the activation energy in the present temperature range should be EA ~ Eo - 4 kcal/mole, where Eo is the threshold energy and is probably less than 30 kcal/mole (perhaps much less). From this, one obtains k -~ 1014 exp l e o - 4 kcal/mole]. This is what we have always used in our modeling. If one applies the same considerations to HCO (very similar to NNH), one gets very close to the Schecker and Jost dissociation rate coefficient expression.

Author's Reply. We did not look for NO2 in these particular experiments because we've looked for it in previous studies and found none. Likewise our

KINETICS AND MECHANISM OF NH3 OXIDATION computer model, which includes all the known HO2 and NO2 chemistry but which does not include the unknown N2H + 02 reaction, predicts no detectable NO~. How including the N2H + 02 reaction would change this I couldn't say. As to your second point, you may be right but please r e m e m b e r that in this paper we're only trying to account for the reaction kinetics at 1300~ K. In an extension of this work which is beyond the scope of this paper we do account for the kinetics over a considerable temperature range. As you might expect in this extension the rate constants at 1300~ K are not altered but the A factors and activation energies are.

105

J. Levy, MIT Energy Laboratory, USA. Have you examined the applicability of your mechanism to higher temperature flame conditions, or, alternatively, would you care to speculate on its applicability? Author's Reply. Tony Dean has gone on to do work on NH 3 flames, both computer modeling and measurement by laser. That's an extension of the model both to high temperatures and high NH3 concentrations. Our current version of the model does reasonably well but there are difficulties which seem to be associated with the high NH3 concentration rather than the high temperature. Apparently for flames we need to include reactions which are second order in NH,.