The mechanism of no formation from nitrogen compounds in hydrogen flames studied by laser fluorescence

Eighteenth Symposium (International) on Combustion

The Combustion Institute, 1981

T H E M E C H A N I S M O F NO F O R M A T I O N FROM N I T R O G E N C O M P O U N D S I N H Y D R O G E N F L A M E S S T U D I E D BY LASER FLUORESCENCE C. MORLEY Shell Research Limited, Thornton Research Centre, P. O. Box 1, Chester CH1 3S H

Laser fluorescence was used to detect OH, NH, CN and NO in rich premixed H 2 / O 2 / A r flames at 1790-2200 K doped with CHaCN. The results are consistent with NO and N 2 being formed by reactions involving nitrogen atoms N + OH---~ NO + H

(21)

N + NO---) N 2 + O

(10)

with k21/klo = 1.0 + 0.3. N atoms are formed by the essentially irreversible reaction NH + H ---) N + H 2

(11)

k H ~ 5 • 10 -H cm a molecule -~ s -~ with the accuracy limited by the calibration for NH. Routes through NH and NH 2 are insignificant in these flames. Approximate upper limits on the reactions of NO with NH and NH 2 are 7 • 10 -l~ and 8 • 10 -~2 cm a molecule -1 s -~ at 1790 K. The results do not rule out the possibility that the route through NH 2 proposed by Fenimore can become significant in flames of low radical concentration.

1. Introduction

themselves and to lead very quickly to NH or NH z.

The chemical mechanism by which nitrogen compounds are interconverted in flames has been subject to several investigations ~-~ but is still not completely understood. The following outline scheme seems to be agreed.

HNCO + H---) NH~ + CO

(3)

NCO + H ~ NH + CO

(4)

The NH~ species (NHa, NH~, NH, N) can be interconverted by reactions of the type NH x + H---) NHx_ 1 + H 2

Fuel nitrogen ---) HCN ---) N H N x~ O /

and one or more of them can react with an oxidizing species, Ox, to form NO or react with NO to form N 29

NO",

N2

The conversion of fuel nitrogen to HCN seems to take place efficiently and the rate limiting step in many flames is the oxidation of HCN. The most important reaction appears ~'a'9 to be: HCN + OH ~ H O C N + H

(5)

(1)

NH~ + Ox ---) NO + ... NHx + NO ---) N 2 + ...

(2)

The cyanate species (HOCN, HNCO and NCO) produced from this and other proposed4'0'9 oxidations of HCN are likely to be interconverted among 23

(6) (7)

Which of NH x species takes part in this dividing point in the mechanism is controversial. All the nitrogen-forming reactions NH 2 + NO --~ N~ + ...

(8)

24

COMBUSTION GENERATED P O L L U T I O N NH + NO ~ N 2 + ...

(9)

N + NO---~ N 2 + ...

(10)

are known to be fast at room temperature. 1~ In flames, Haynes 2 presents evidence for both N 2 and NO being formed from N atoms, whereas Fenimore ~ prefers NH 2. De Soete 7 favours N atoms, Seery and Zabielski ~3 favour NH, and Duxbury and Pratt's" modeling of shock tube experiments required both NH and NH 2. It is probable that one of the causes of these divergent views is that the dominant mechanism depends on the conditions. In H2/Oz/Ar flames near 1800 K we find in this study that, in agreement with Haynes, 2 reactions (6) and (7) involve the nitrogen atom but, contrary to his assumption, that it is not equilibrated with the rest of the NHx species and that the limiting step is NH + H---~ N + H 2

(11)

NO formation has been analysed by others in terms of behavior equivalent to the oxidizing species Ox in reaction (6) being 02, 7 OH, s or H20. ~ In our flames we find that it is OH. In hydrocarbon flames there is a process 3'4 which regenerates HCN by reaction of hydrocarbon fragments with NH x and with NO. This complicates the interpretation of experimental results in hydrocarbon flames. Because the conditions governing the removal of hydrocarbons and HCN are broadly similar there is the likelihood that the H C N decay, and hence much of the subsequent chemistry, takes place in regions of the flame where hydrocarbons are still present. In order to study the processes leading to NO and N 2 formation without this interference we have used H J O 2 / A r flames. These also have the advantage of having a wide variation of radical concentration within a particular flame, allowing a more stringent test of mechanisms.

2. Concentration Measurement by Laser Fluorescence In most previous studies ~-3'5-sconcentrations have been measured using probe sampling. Although the assumption that HCN, NH 3 and NO are efficiently sampled from the reactive conditions of flames at atmospheric pressure has been made, it does not seem to have been comprehensively tested. We have avoided these difficulties by using the technique of laser fluorescence to measure the concentration of certain species in situ. This aspect of the work will be described in more detail elsewhere. 9 Laser fluorescence has been demonstrated for molecules in flames by several groups 15-~9but, apart from Ref. 17, there has been little quantitative use of derived concentrations. One of the reasons for

this is the occurrence of fluorescence quenching which makes the conversion of fluorescence signals to concentrations difficult. In premixed flames the major quenching species have constant concentration throughout the useful part of the flame so that the signal is proportional to the concentration of the fluorescing species in each flame. For measurements of CN this proportionality is sufficient because its behavior is kinetically first order and only relative concentrations within a single flame are required. For OH the proportionality constant for each flame was calculated from the concentrations of the major species and their quenching rates. For the unsaturated conditions used [OH] = S c k.2 o IH20] 9 (1 +

X(k,/k,2o)([M,] / [H20])}

where S is the fluorescence signal, [M,] and k, are the concentrations of the major species and their quenching rates and c is a constant including laser intensity, sensitivities, geometrical factors, etc.; the summation is over major species other than water. Measurements 9 in a series of flames of known [OH] yielded the relative quenching rates kJkH2o and the term c k H o, allowing [OH] to be calculated in any flame, izi was found that k . o:kH :kco = 1.0 : 0.32 + 0.10 : 1.27 + 0.20 independent of t~mperature. Our OH measurement are probably accurate to within 25%. In hydrogen flames the NO signal could be calibrated by direct addition of NO to the flame but separate quenching measurements showed that the quenching was almost entirely by H20. NH concentrations in rich ammonia flames have been measured by Fisher z~ using optical absorption. We have obtained a calibration for NH fluorescence, using these flames and the assumption that water is the only quecher, but the difficulty in reproducing the conditions using different burners means that the calibration can only be considered a rough one. Although fluorescence from the visible system of NH 2 has been readily observed by others 2~ at low pressures, we were unfortunately unable to detect it in any kind of atmospheric pressure flame. The weakness of the transition, to which the fluorescence is particularly sensitive in heavily quenched environments, is probably responsible.

3. Experimental Flat premixed flames at atmospheric pressure were burnt on a burner made from a bundle of tubes. 4 For those with a burnt gas velocity less than 5 m s -I the burner diameter was 20 mm, for those greater it was 10 mm. A shield flow of argon on the outside reduced air entrainment. Some of the flames were

T H E MECHANISM OF NO FORMATION IN HYDROGEN FLAMES from isothermal sets previously characterized; 2~ some sodium line reversal temperature measurements were made; some flame temperatures were interpolated using these data. Characteristics of the flames used here are given in Table I. The laser fluorescence apparatus will be described 'elsewhere.9 Briefly, a nitrogen laser (Molectron 300 kW) pumped a dye laser (Molectron DL100 + a telescope) which gave a 10 ns pulse of several tens of microjoules between 380 and 680 nm. The output could be doubled by KDP crystals to give 0.5 to 3 ~tJ between 280 and 340 nm or by a KPB crystal to give <0.1 ~J at 226 nm. The incident beam was focused to <0.5 mm diameter in the flame, and the fluorescent light was filtered by appropriate glass and liquid filters or a small monochromator. The photomultiplier signal was gated for 500 ns during and after the laser pulse. Table II shows the incident and fluorescent wavelengths used to detect each species. JANAF thermochemical data zz have been used.

4. Results and Discussion HCN Decay Figure 1 shows the CN, NH and OH fluorescence signals in flame B to which a small quantity (<0.1%)

[HCN]/IO

25

-4

F-

[CN]/5~ 0 -~ 0

0

r

i

i

4

8

12

DISTANCE,mm

F1c. 1. Concentration of various species in flame B containing the equivalent of 17.8 x 10 -4 mole fraction of CHaCN. The CN, OH and NH profiles are from fluorescence measurements; the HCN profile is derived from those of CN and OH. The HCN and CN are put on an absolute basis assuming 100% conversion of CH3CH to HCN in the reaction zone. The calibration for NH is from Fisher's z3 ammonia flames and is only approximate.

TABLE I Characteristic of flames Burnt gas velocity

Average temperature Flame

Hz/ Oz/ Ar

K

A

5.22/1.0/6.03

B

4.44/1.0/6.56 2.99/1.0/8.43 5.00/1.0/4.26 3.91/1.0/3.04 3.15/1.0/4.85

1790 1790 1790 2000 2130 2200

C D E F

ms

8.6 6.9 6.3 11.8 20.8 28.6

Measured k2

Measured k~l

35}

5.0

3.3 7.6 6.7 9.7 12.5

5.0 69 7.6

a = 10 -~a cm a molecule -1 s -1 b = 10 -ll cm a molecule - l s -1

TABLE II Species observed by laser fluorescence

OH CN NH NO

Fluorescence

Excitation

Species A-X, B-X, A-X, A-X,

(0,0) (0,0) (0-0) (0,0)

Pl8 bandhead Q branch Pll bandhead

285.6 388.5 336.0 226.4

nm nm nm nm

(0,0), (0,1) (0,0) (0,0) (0,1)-(0,5)

300-320 380-390 336 235-280

nm nm nm nm

26

COMBUSTION GENERATED POLLUTION

of CHaCN had been added. The OH signal has been converted to a concentration using our quenching data. The HCN profile can be deduced from those of OH and CN by use of the balanced reaction: HCN + O H ~ CN + H20

(12)

It can be seen in Figure 2 that the [OH] profile passes approximately through the maximum in the [NO] profiles. This is what would be expected 2 if Ox were OH, and NH, species were the same in reactions (6) and (7), and the rates of (6) and (7) were equal. The equations are then:

A systematic study of the kinetics of HCN decay will be reported elsewhere 9 but within a particular flame it can be represented by HCN + OH ~ Products

(2)

NHx + OH---) NO + ... NHx + NO

[CN] [OH]

~ Na + ...

(7)

We have modeled the system with these two equations and with the NH x coming from HCN oxidation

From (12) and (2) it follows that: in--

(15)

HCN + OH ~ NH x + ...

(2)

t

= -k2

i

[OH] dt + constant (13) o

Assuming NHx to be in a steady state results in the rate equations:

Plots of the left hand side of eq. (13) against I OH dt gave good straight lines and the derived values of k 2 are given in Table I. CN signals were proportional to the CHaCN for additions up to about 2 x 10 -3 mole fraction. Above this, OH concentrations began to be reduced and the HCN decay slowed.

d [NO]

--=

k 2 [HCN]

dt

[OH]

9 .( [OH][OH] + (kr/k,,) ~ - [NOl ) d [HCNI dt

k, [HCN] [OH]

(16)

(17)

NO Formation A calibration for the NO fluorescence signal was obtained by adding NO alone to the flame. The NO signal was found to be constant through the burnt gas and proportional to the amount of NO added. The constancy suggests the absence of large quantities of NH, species in the burnt gas, otherwise NO would be removed by reaction (7), and the linearity suggests that an insignificant amount of N 2 is formed. The most well known reaction for removing NO

The NO concentrations shown by lines in Figure 2 were calculated by a Runge-Kutta integration of (16) and (17) using the following data. OH concentrations were from measurements but for ease of computation were fitted (quite closely) to a sec-

8 \\

:

IO~[HCN]O 9,o

3

6

H + N O ~ N + O H A H = 205 kJ mole -~ (14) is too slow ~2 to destroy appreciable amounts under these conditions. We therefore assume no loss of NO in the calibration. The NO concentrations (see Figure 2) in flames with added CH3CN showed highly non-linear behavior and there is a maximum in the NO profile when the additions are large. This means that the NO-producing reaction (6) is initially faster than the NO-consuming reaction (7) but later in the flame becomes slower. This could have two causes. Either the NH~ species in each reaction are different and that in reaction (6) decreases relative to that in (7) down the flame, e.g. by its value of x being lower, Or the concentration of the oxidizing species, Ox, in (6) becomes less down the flame. Neither of these cases is satisfied by Fenlmore " ' s proposal that the rate of the oxidation is proportional to [NH~] [H20 ] .

0

4

v.o

-

[o.~'---

/~

z~ 2

l

o o

4

~

1'2

DISTANCE,mm

FIG. 2. NO profiles in flame B. Points are measured, lines are calculated from the measured CN and OH concentrations using kT/kls = 1.0, see text.

T H E MECHANISM O F NO FORMATION IN HYDROGEN FLAMES ond-order decay, k 2 was from the analysis of the CN decay described earlier, kT/k15 was taken as 1.0 and quantitative conversion of CH3CN to HCN at time zero (the reaction zone) was assumed. The calculation reproduces all the qualitative features of the NO profile including the maxima at large additions. It also simulates the interesting inverse dependence of NO on fuel nitrogen addition which occurs in the later part of the flame at large additions of CHaCN. Fits of similar quality were obtained in all the flames in Table I using the same value of kT/ktn, and Table III shows that absolute yields of NO are fitted well. The calculated concentration of NO late in the flame were approximately proportional to the assumed ratio kts/k 7 at high fuel nitrogen additions but less sensitive to it at low additions. Any inaccuracy in the OH

27

calibration would be reflected directly in kT/kl~. The overall uncertainty in the fitted value of kT/k~5 is estimated to be 30%. When the CH3CN additions was replaced by similar quantities of NH 3 larger amounts of NO were produced and the NO profiles were fiat throughout the burnt gases of the flame. This implies that NH 3 reacts very quickly in the reaction zone and this fast reaction of presumably all the NH~ species supports the steady-state assumption in the above simulation. The larger amounts of NO formed from NH 3 compared with CHaCN reflect the higher concentrations of OH present while the NH~ species decay. However, it was not possible to model the amount of NO formed from various additions of NH 3 using reactions (7) and (15) with kT/k15 = 1.0 and a constant OH concentration. Presumably the

TABLE III Observed and calculated NO concentrations [NO] at 12 mm Flame

Fuel nitrogen

Observed

Calculated

7.3 14.4 25.2 43.2

2.4 3.3 3.8 3.6

2.5 3.3 3.6 3.2

9.0 17.8 31.0 53.0

3.2 4.3 4.9 4.4

3.1 4.2 4.7 4.5

31.0 'NHz 53.0 9

4.9 7.9 10.3 12.3

B

9.9 19.6 34.4 58.6

9.0 10.8 9.6

6.1 9.0 11.4 9.0

D

5.9 11.6 20.4 34.8

3.4 5.6 8.8 8.1

3.4 5.3 6.9 8.1

E

21.2

16.0

14.8

F

15.9

I 1.5

11.3

6.4

84

Concentration are in units of 10-4 mole fraction and for the fuel nitrogen are the equivalent concentration in the burnt gas. Except where noted the fuel nitrogen was CH3CN.

28

COMBUSTION GENERATED P O L L U T I O N

NH 3 reacts while the OH concentration is still rising, or, under the high [NH,] conditions present in and near the reaction zone, it is conceivable that NH x species can react among themselves to form N 2.~

Nature of NH, Species Forming NO and Ne The modelling above did not require knowledge of which NH x species was involved in the branching step. To obtain this information measurements on one or more of the NHx species is required. We have observed NH by fluorescence but have only a rough calibration at present (q.v.). This allows us to say something about the mechanism but not to measure accurate rate constants. In all hydrogen flames doped with CH3CN the NH signal decreased monotonically down the flame and was proportional to the added CH3CN throughout. When the additive was NH 3 no NH signal was observed and this is consistent with the constancy of the NO concentration under the same conditions. The crucial observation was that the NH concentrations in flames doped with CH3CN were unaffected when NO was also added. NO concentrations in flame B at 2.5 mm with [HCN]o = 17.8 • 10 -4 could be increased by a factor of 3 with a less than 10% change in NH signal. Similar behaviour was observed in all the flames in Table I. This implies that NH, or any other NH, species equilibrated with it, does not react appreciably with NO under these conditions. The linearity of the NH with C H 3 C N addition also supports this conclusion, indicating as it does the absence of a process involving NH other than first order in nitrogen compounds. Of the possible N 2 forming reactions: NHe + NO ---) N2 + ...

(8)

NH + NO ---) N2 + ...

(9)

N + NO ---) N 2 + O

(10)

reaction (9) is excluded by the evidence above. Whether it is consistent with (8) being the important reaction depends on which NH x species is directly produced from the HCN oxidation. If it is NH 2 (perhaps via reaction (3)) then (8) together with its associated oxidation reaction: NH 2 + OH---) NO + ...

(15)

cannot be a major route for NHx removal. This is because NH would be formed from NH 2 via NH, + H ~ NH + H 2

(18)

NH 2 + OH ~ NH + H20

(19)

and so its concentration would depend on [NH~]

which in turn would depend on [NO]. If H C N oxidation produces NH then (8) and (15) can be important only if their combined rate is sufficiently large for [NH2] to be strongly depleted, so making reactions (18) and (19) occur effectively only in the reverse direction. The rates of these reactions are unfortunately not known. It has been found z3 that NH and NH~ are equilibrated in some ammonia flames but conditions in them are sufficiently different to make it unsafe to extrapolate this conclusion to our hydrogen flames. It is unfortunate that we were unable to detect NH 2 directly because its observation would provide direct evidence on the occurrence of reaction (8). Turning to the other possible route via nitrogen atoms, we show now that the formation of N atoms from NH is essentially irreversible under our conditions: NH + H--~ N + Hz

AH = - 1 2 3 kJ m o l e - '

NH + OH---) N + H20 AH = - 1 8 8 kJ m o l e - '

(11)

(20)

For this to occur the rate of (10) together with (21) N + OH ---) NO + H

(21)

have to be very much faster than reverses of (11) and (20). The latter is too endothermic to contribute significantly so that k,o [NO] + k2, [OH] > > k,, [H2] With [OH] = [NO] = 5 • 10 -4 , [H:] = 0.22 and klo = k21 = 5 • 10 -H cm 3 molecule -l s -l (flame B) one obtains kl~ < < 2 x 10 -~3 cm ~ molecule -1 s -~ The endothermicity means that this is certainly true, since the exponential factor is smaller than 3 • 10 -4 at 1790 K. It is likely that N atoms will be below their partial equilibrium level in most flames except those with very low concentrations of both NO and OH. This non-equilibration of N with NH means that reaction (10) being the nitrogen-forming step is consistent with NO not affecting the NH signal. The above analysis means that in these flames the route to NO and N 2 formation passes through NH and that the reaction removing NH is essentially irreversible. We can therefore write d([HCN] + E[NH.])

dt

- r [NH]

(22)

where E[NHx] = [NH~] + [NH2] + [NH] + [N] and r is the removal rate for [NH]. In the modelling

THE MECHANISM OF NO FORMATION IN HYDROGEN FLAMES of the NO profiles we effectively assumed that dZ [ N H , ] / d t was small compared with d [ H C N ] / d t and its success helps to validate the assumption; additional evidence is presented later. Equation (22) then becomes: d [HCN]

r [NH]

dt

k,5 [OH] + k 7 [NO] d [NO] - = r[NH] kis [OH] - k 7 [NO] dt

(25)

where only the NH x species taking part in reactions (7) and (15) is assumed to be in a steady state. With k~5/k7 = 1.0, as found earlier, this becomes

(23) 1

[OH + [NO]

d[NO] -

The left hand side can be obtained from the empirical equation (17), giving k~ [HCNI [OH]

29

=r

(24)

[NH] In Figure 3 the left hand side of (24) is plotted against [H] obtained from the measured [OH], taking account of the temperature reduction caused by radical excess and calculated assuming that the flame is adiabatic in the burnt gases. An alternative way of obtaining an expression for r uses measurements of nitric oxide. It can be shown that reactions (7) and (15) (with Ox = OH) lead to

FLAME

0

EQU24

EQU26

A

1.50

Z~

9

B

2.22

0

9

C

2.61

[]

[]

[NH] [OH] - [NO]

-

= r

(26)

dt

This method of obtaining r is less accurate than using (24) because of the subtraction of comparably sized quantities, [OH] and [NO]. Results were in fact obtained with [OH] both greater and less than [NO]. Some of the results are shown as closed circles in Figure 3 together with their estimated errors. The agreement with the other method is satisfactory. Figure 3 shows that r is proportional to [H]. This holds for a range of stoicheiometries, whereas a similar correlation with [OH] is not found. Hence NH is being removed largely by reaction (11). This produces nitrogen atoms and therefore N~ and NO are mainly formed by the N atom reactions (10) and (21). Had NH 2 been responsible NH would have

O

i l.,9

s 0

2

, ~ / 0

kJ

_1 5

!

I

I

I0

15

20

[H ] , 10 -3 mole fraction FIG. 3. The rate of removal of NH, r, as a function of H atom concentration in three H J O J A r flames at 1790 K. r is obtained from either equation (24) or equation (26), see text, and its absolute value is only approximate.

30

COMBUSTION GENERATED P O L L U T I O N

been removed by (18) and (19) and r would have been independent of the radical concentration. There is no apparent intercept on the line in Figure 3, suggesting that at least 80% of flux through NH proceeds to N atoms. The approximate calibration from Fisher's ammonia flames ~ has been used to put the values [NH] in Figure I and r in Figure 3 on an absolute basis. The NH mole fraction in Figure 1 is about 2.5 • 10 -7 (40 times smaller than the average in Fisher's flames). If NH 3, NH 2 and NH are assumed equilibrated then their total mole fraction is 7 • 10 -8, a factor of 70 less than HCN or NO, justifying the neglect of the E[NHx] term in equation (22). Using the approximate calibration for NH the slope of the line in Figure 3 yields kit = 5 • 10 -H cm 3 molecule-is -~ at 1790 K, accurate to perhaps a factor of 2. Similar analysis of the less extensive results from the higher temperature flames leads to values of k ll shown in Table I. There is a tendency for k H to increase with temperature although there are insufficient data to say whether this trend is real. NH signals tended to be smaller in higher temperature flames because of faster removal and fluorescence quenching by the increased of concentrations of H atoms and H~O respectively. An upper limit can be put on the reaction of NH with NO. The insensitivity of [NH] to added NO means that

which under the most restrictive conditions we studied (flame C at 14 mm, < 10% change for A [NO] = 1.3 x 10 -3, [HI = 1.8 x 10 -s) leads to k9 < 7 • 10-m cm-3 molecule-i s -1 A room temperature rate for this reaction H is 4.7 • 1 0 - " cm 3 molecule-' s -1, implying a negative temperature dependence or alternatively products which quickly regenerate NH x species. Reaction (8) seem to be unimportant in these flames, which means that the flux through it is small (<20%) compared with the flux through NH, i.e. (27)

If NH 2 is assumed equilibrated with NH, conditions in all the flames at 1790 K lead to ks < 8 • I0 -'~ cm 3 molecule-'s- l

Haynes 2 interpreted the NO formation in his ethylene flames by the same general mechanism-reactions (7) and (15)-- and obtained k~5/k 7 = 1.0, in agreement with this work. He also concluded that the nitrogen atom was the species involved in the branching steps but he assumed that N was equilibrated with the rest of the NH x pool. As discussed earlier and as pointed out by Fenimore l this is unlikely to be the case. However, this would cause fairly small errors in his deduction of the mechanism, which relied essentially on finding the order with respect to radicals of NH 3 disappearance. His assumption that the rate limiting step involved N atoms, i.e. NH 3 removal rate a [N] a [NH3] [radicals] 3 is nearly equivalent to ours with reaction (11) rate limiting: NH 3 removal rate a [NH] [HI a [NH3] [radicals] ~ [HI if NH and NH 3 are equilibrated. It can be shown that the value Haynes deduced for k'm, is related to our ka, by

k'lo

kg[NO] < < kll[H]

ks[NHz] [NOI < < kii [NH] [H]

Comparison with Other Work

(28)

The room temperature rate is 2 • 10 -H cm 3 molecule -~ s -1 with a negative activation energy of about 6 kJ mole-~. 1~ If this dependence is continued up to 1790 K the rate would be 2 • I0 -12 cm ~ molecule -1 s -l, which fulfills (28). Thus the unimportance of reaction (8) in our flames is consistent with the low temperature rate constant.

[H~I

-- = kl, K l l ( l O H ] + [NO])

(29)

where Kll is the equilibrium constant of reaction (11). As noted by Haynes himself, 2 his value of k'lo increased with the richness of the flame when [ H J increases and ([OH] + [NO]) decreases, in accordance with (29). For his ethylene/air flame at 2070 K, ~b = 1.52, using [ H J = 5.7 • 10 -2 , K~I = 310 and his measured values OH = 5 • 10 -4, NO = 3 • 10 -4 and kio = 1.7 • 10 - u cm 3 molecule-' s -~ gives kit = 8 X 1 0 - " cm a molecule -1 s -t in agreement with our rough value of 5 • 10 - u cm a molecule -I s -1. It should be pointed out that Fenimore 1 was able to show that Haynes' results 2 are consistent with his mechanism involving NHa (q.v.). In addition there is a possibility of interference from hydrocarbons in these ethylene flames making the interpretation liable to error. Fenimore 1 found in H J N a O flames that N~ formation was by reaction (8) with a rate of 8 • 10 -12 cm 3 molecule -1 s -~. Such a value is just about compatible with our results expressed as the inequality (28). These flames have relatively low radical concentrations (near equilibrium) and it may be that the dominant mechanism changes. In flames with low radical concentration (H2/N~O flames) the NHa route may be most important, whereas under high radical conditions ( H J O ~ flames) the route

T H E MECHANISM OF NO FORMATION IN HYDROGEN FLAMES

31

6. B. S. HAYNES,D. IVERACHAND N. Y. Kmov, Fifteenth Symposium (International) on Combustion, 1103 (1975). 7. G.G. DE SOErE, Fifteenth Symposium (International) on Combustion, 1093 (1975). 8. C. P. FENIMORE,Combustion and flame, 19, 289 (1972). 9. C. MORLEY,to be published. 10. W. HACK,H. SCHACKE,M. SCHROTER,H. WAGNER, Seventeenth Symposium (International) o n Combustion, (1979). Ii. I. HANSEN, K. H()INGHAUS, C. ZETZSCH AND F. STUHL, Chem. Phys. Letters, 42, 370 (1976). 12. D. L. BAULCH, D. D. DRYSDALE, D. G. HORNE, Evaluated kinetic data for high temperature reaction, Vol. 2, Butterworths, London, 1973.

through N atoms, which depends more strongly on radical concentration, may become dominant. In agreement with this, preliminary measurements of NH in H J N 2 0 flames show that, in contrast to H 2 / O ~ flames, its concentration is affected by additional NO and hence that routes through NH and NH~ may be significant. The C2H~/O. flames of Haynes ~ are intermediate in radical concentration and both mechanisms have been shown to be roughly compatible.

5. Conclusions The route for interconversion of nitrogen compounds in rich H s / O J A r flames at 1790-2200 K is

NO CHaC N

100% ,o

I~ HCN O - .HQ ~ (NH~ kl,[H]. N

r:y:o

~

''f

L"

[ O1

N~ cule-1 s-l. Nitrogen atoms are not equilibrated with the rest of the NHx species. Upper limits on the rates of NH and NH2 with NO at 1790 K are 7 • 10 -Is and 8 • 10 -12 cm 3 molecule-' s - ' respectively. The rate of the latter is compatible with the possibility that a route to NO and N~ becomes important in flames of lower radical concentration.

REFERENCES 1. C. P. FENIMORE, Seventeenth Symposium (International) on Combustion 661 (1979). 2. B. S. HAYNES, Combustion and flame, 28, 113 (1977). 3. B, S. HAYNES, Combustion and flame, 28, 81 (1976). 4. C. MORLEY, Combustion and flame, 27, 189 (1976). 5. C. P. FENIMORE,Combustion and flame, 26, 249 (1976).

13. D. J. SEERY, M. F. ZARIELSKI,Combustion and flame, 28, 93 (1977). 14. J. DvxmJRY, N. H. PRATT,Fifteenth Symposium (International) on Combustion, 853 (1975). 15. P. A. BONCZYK,J. A. SHIRLEY, Combustion and flame, 34, 253 (1979). 16. J. H. BECHTEL,Applied Optics, 18, 2100 (1979). 17. C. H. MULLER, K. SCHOFIELO,M. STEINBERG,H. P. BROIDA, Seventeenth Symposium (International) on Combustion, 867 (1979). 18. M. MAIL)/NDER,J. Appl. Phys., 49, 1256 (1978). 19. H. HARAGUCm,S. J. WEEKS, J. D. WINEFORDNER, Speetro. Chem. Acta A, 35, 391 (1979). 20. G. HANCOCK, W. LANGE, M. LENZl, K. H. WELGE, Chem. Phys. Lett.,33, 168 (1975). 21. C. J. HALSTEAD, D. R. JENKINS, Twelfth Symposium (International)on Combustion, 979 (1969). 22, J A N A F Thermochemical Tables, 2nd Edition, (1971). 1974 Supplement J. Phys. Chem. Ref. Data, 3, 311 (1974). 23. C. J. FISHER, Combustion and flame 30, 143 (1977).

COMMENTS Peter Schug, DFVLR Institute fur Physikalisone, Germany. We have been studying premixed C a H s / O J N 2 and C 3 H s / O J A r flames doped with

CHzNH ~ in the equivalence ratio range 0.8 < ~b < 1.8. Preliminary evaluation of the exhaust gas composition data lead us to conclude that reactions

3~

COMBUSTION GENERATED POLLUTION

o f t h e type

[The s e c o n d part of the q u e s t i o n is a n s w e r e d directly in the paper.] NH~ + NH~-", N 2 + ...

and/or NH~ + N H ~

N2H + ...

N~H ~ N 2 + H c o n t r i b u t e s i g n i f i c a n t l y to the f o r m a t i o n of m o l e c u las n i t r o g e n d u r i n g the c o n v e r s i o n p r o c e s s o f the d o p i n g material to N O or other fixed n i t r o g e n species. C o u l d the fast depletion of N H , t h r o u g h one or both o f t h e s e reaction c h a n n e l s be r e s p o n s i b l e for the low N H c o n c e n t r a t i o n y o u m e a s u r e d ? A s s u m i n g partial e q u i l i b r i u m for the NH~ pool N is the least a b u n d a n t species u n d e r fuel rich conditions. If n o t b e e n equilibrated I w o u l d expect the N a t o m s to be e v e n lower in c o n c e n t r a t i o n d u e to their d e p l e t i o n via reactions (21) a n d (10) (accordi n g to y o u r s c h e m e t h e only outlet for t h e N H , pool). Since y o u f o u n d reactions via N H 2 a n d N H to be insignificant, h o w m u c h slower w o u l d y o u e s t i m a t e their reaction rates to be with respect to reactions (21) a n d (10)?

Author's Reply. I n the flames s t u d i e d in t h i s work the fast r e m o v a l o f NH~ b y N H + H (reaction 11) together w i t h its relatively slow formation from H C N (reaction 2) were s u f f i c i e n t to explain the concentrations of N H observed. Reactions a m o n g the NH~ were i n s i g n i f i c a n t . F o r example, if N 2 were b e i n g f o r m e d b y reaction of N H 2 + N H 2 w i t h a rate constant of 5 • i 0 - H c m 3 m o l e c u l e -~ s - t the r e m o v a l rate o f NH~ in flame B w o u l d be o f t h e order of 4 • 10 t5 m o l e c u l e c m -3 s -~. R e m o v a l b y reaction (11) is a b o u t 10 ~ m o l e c u l e cm -~ s -t. In f l a m e s w i t h lower quantities o f radicals p r e s e n t it m a y be that the N H x concentration can b u i l d up, since, as d e s c r i b e d in the paper, NH~ d e s t r u c t i o n d e p e n d s on [radicals] 3 w h e r e a s its f o r m a t i o n from H C N d e p e n d s o n l y o n [radicals]. R e a c t i o n s a m o n g the NH~ s p e c i e s m a y t h e n b e c o m e a p o s s i b i l i t y a n d I am interested that y o u have e v i d e n c e for their occurrence. As m e n t i o n e d briefly in the paper, they m a y also play a role u n d e r high radical c o n d i t i o n s w h e n N H a is a d d e d to h y d r o g e n flames.

1. Wolfrum, MPI fur Stroemunsforschung, West Germany. W i t h respect to the p r o d u c t s of t h e reactions (8) a n d (9), I w o u l d like to p o i n t o u t that beside the direct formation of m o l e c u l a r nitrogen, other p r i m a r y p r o d u c t s are possible, i.e. N H 2 + N O --~ N~H + O H TM a n d N H + N O ~ N 2 0 + H I~ w h i c h s h o u l d be i n c l u d e d in a proper kinetic m o d elling.

REFERENCES [1] A. JACOBS, O1PLOM. Thesis, Universit~it, G6ttingen, 1980. [2] M. GEARING, K. HOYERMANN,H. G. WAGNER AND J, WOLFRUM BER. Bunsenjes, Physik. C h e m . , 73, 956 (1969), 75, 1281 (1971).

Author's Reply. T h e p r o d u c t s N2H a n d N 2 0 w o u l d p r o b a b l y be c o n v e r t e d rather q u i c k l y to N 2 a n d the detailed course of the reaction w o u l d n o t affect o u r c o n c l u s i o n on the relative u n i m p o r t a n c e of t h e s e reactions in o u r flames.

John W. Daily, University of California, USA. W h a t line d i d y o u p u m p in O H ? We have s h o w n that in the weak excitation limit one s h o u l d p u m p near the peak of the rotational distribution. T h i s m i n i m i z e s the t e m p e r a t u r e d e p e n d e n c e of fluorescence signal o n O H concentration.

Author's Reply. T h e P,8 line of the (1,0) b a n d was used. A P b r a n c h line w a s preferred to o n e from the s t r o n g e r Q b r a n c h in order to m i n i m i z e the p o s s i b i l i t y of s i g n i f i c a n t absorption of the incid e n t light before it r e a c h e d the scattering point, a n d the particular line w a s c h o s e n for t h e reason y o u stated.