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The pyrolysis and oxidation of hydrazine behind shock waves
Tenth Symposium (International) on Combustion,
pp. 353-364,
The Combustion'Institute, 1965-
THE PYROLYSIS AND OXIDATION OF H Y D R A Z I N E B E H I N D S H O C K WAVES K. W. MICHEL, AND H. GG. WAGNER
Institut fiir physikalische Chemie der Universit~t GOltingen, G6ttingen, Germany The thermal decomposition of hydrazinc highly diluted with Ar or He (0.03%-0.5% N2H4) has been investigated spectrophotometricaUy behind shock waves in the temperature range of 1100~ to 1600~ At the lower temperatures and higher concentrations of N2It~, the absorption traces recordcd on the oscilloscope exhibit an induction period and a certain dependence of the half-lives of N~H4 on concentration, viz., the characteristics of a chain decomposition. As one approaches higher temperatures (T > 1300~ and lower partial densities of N_~H4, the rate law for the decomposition becomes roughly first order; the first-order rate constant being independent of N2H4 concentration but revealing a slight dependence on total gas density pv: log k ~ 12.8 - [-(52 kcal/mole)/2.3RT] at or = 7.5 X 10-s inole/cm~ and log k ~ 12.0 - [(48 kcal/mole)/2.3RT] at pT = 2.5 X 10-6 mole/era 3 (k in see-0. The addition of oxygen to an 0.3% N2H~-Ar mixture causes noticeablc acceleration of the decomposition only when added in appreciable amounts ([-O2-]/[N~H,-] > 2). The amount of NHs formed during the pyrolysis of N2H4 has been de~rmined by a combined incident-reflected shock-wave technique, again using uv absorption at 2300 A for quantitative determination of both reactant and product. The ratio rNHz-](formed)/[-N;H~-](decomposed) decreases from 1 at 1100~ to 0.5 at 1600~ and is about 0 at 2000~ At temperatures above 1400~ the NH radical could be detected photometrically. Introduction The investigation of the thermal decomposition of gaseous hydrazine by classical methods suffers greatly from wall catalysis. 1,2 To obviate this difficulty, other techniques have been resorted to such as decomposition flames, 8 flash photolysis, 4 studies of explosion characteristics. 5 Information about the controlling processes under nearly isothermal conditions, however, may be obtained most conveniently by means of shockwave techniques. The single-pulse method has been applied to the N2H4 pyrolysis by Moberly e and by McHale et al. 7 in the temperature range 930~176 and 950~176 respectively. The data of the latter investigation has been interpreted in terms of a first-order rate law, giving k = 1.1 X 10~3 exp ( - 5 5 kcal/RT)sec -1. Mass spectroscopy has been combined with the shock-wave technique by Dicsen s in a study of the N2H4 pyrolysis and at temperatures from 1200~ to 2500~ at appreciably lower pressures
(total gas density, pr = 7 X 10-7 to 13 X 10-~ mole/era 3) than in the above-mentioned investigations and in this one (PT = 2.5 X 10- ~ to 14 X 10-5 mole/cm3). In this study a uv-absorption technique was used, permitting the continuous recording of the kinetic behavior of hydrazinc, which was present in comparatively low concentrations (less than 0.5% N2H4 in Ar or He). Supplementary information could be derived by tracing the concentration course of the NIt radical and by determining the amount of NHs produced in tile N~I/4 pyrolysis. The primary goal in the oxidation experiments was to test the effect of small amounts of oxygen upon the pyrolysis. Apparatus and Method Three different shock tubes, each made of aluminum, wcrc used, in order to eliminate the effects of special properties of the apparatus (boundary-layer effects, attenuation, catalytic
353
354
REACTION
KINETICS
Fro. 1. N2H4 pyrolysis observed at 2300 .~. M = 2.11; re~u1 = 0.596; ,Au/u~ = 3.0%/m; 0.28% N2H4 in Ar; sweep 200 t~see/cm; p5 = 24 )< 10s mole/cm3; pT = 8.6 )< 10-s mole/cmS; Ts = l l20~ decomposition of hydrazine prior to shock-wave experiment, etc.) upon the results. Two square shock tubes, 3.2 cm i.d., one being used without any further treatment and the other having been eloxized, each consisting of a 180-cm-long highpressure section and a 250-cm-long test section. Thus, reaction times behind reflected shocks of at least 1 msec were possible (longer ones were not considered suitable because of deviations from ideal shock behavior such as cooling from the walls and boundary layer effects) ~e. Ten cm from the reflecting endplate, quartz windows (Ultrasil, 90% transmission at 2300 ~,) were inserted flush with the walls for spectrophotometric recording with uv light. The third tube with a round cross section of 10 cm i.d. had a 280-cm-long high-pressure section and a 420-cm-long test section. Quartz windows were inserted 5 cm from the closed end. The leak rate of the low-pressure section was not more than 10--5 mm IIg/min. The tubes were equipped with separate intake and pumping valves, so that the reaction mixture, prepared in a glass system with Teflon joints, could be flushed through the tube. for a length of time before each experiment. A schlieren set-up was used for measuring the speed and attenuation of the incident shock waves, such that the instantaneous shock velocity was measured to within 0.2%. The shock-wave attenuation in the large tube varied between 0 and 1.6%/m, in the small one it was about 3.6%/m for both carrier gases, Ar and tie. Reflected shock speeds could be determined to 1.5% accuracy and provided, together with the attenuation values and the correspondence oi the measured Mach Number with the initial pressure ratio of driver and test gas (within 30% of theoretical), a convenient control for the stationary character of the shock waves. An effective M a t h Number had been defined to account for deviations from ideal shock
r] = 445 gsec.
behavior and proved to be consistent when comparing results from different tubes and from shocks with different attenuation. An almost parallel beam of ultraviolet light from a Xe high-pressure arc (Osram, XBO 150 W) passed through the quartz windows near the end of the test scctioa and onto the slit of a monochromator (Zciss M4 Q II; an EMI-6256 A photomultiplier was used as detector). The incident-light intensity was measured before and after each run with a rotating scctor of 2.8 kcps. The electrical time constant of the recording lmit amounted to 0.4 or 1 ttsec. Measurements of the N2tI4 and NHa concentrations were made at 2300 and 2500 ~, chosen so that the maximum absorption always fell into the most favorable range of 10% to 70%. The absorption traces were evaluated with an accuracy of 5%. Figure 1 shows a typical oscilloscope record of the absorption of hydrazine. The first step in the absorption trace occurred with the passage of the incident-shock front. The trace then remained constant for some 400 t~sec, indicating that there was no appreciable decomposition of hydrazine within 1 msec (particle time scale) at 700~ Even behind the reflected shock, fast decomposition of hydrazine did not occur before a certain induction period had elapsed. The data for characterizing an experiment are given in the legend: M is the effective Mach Number of the incident shock, A u / u its attenuation per m, pT the total gas density at reaction conditions, T5 the calculated temperature behind the reflected shock, and r89 the measured half-life of hydrazine. Further experimental details have been described elsewhere. ~n'13
Reactants For preparing N2H4-Ar mixtures, part of the Ar-stream was bubbled (at known pressure)
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
through liquid hydrazine in a double saturator at 32~ Because N2H4 evaporates sluggishly, this was not enough for accurately establishing the concentration, and the gas laden with N2H4 vapor passed through a condenser, the temperature of which was kept constant at 22.0 ~ -40.05~ Beyond the condenser, the Ar stream saturated with hydrazine was reunited with the main stream of pure Ar or Ar-02 mixtures. The ratio of flow velocities was fixed by a set of properly gauged capillaries or by Bodenstein valves. N2H4 is strongly adsorbed by the acidic aluminum oxide at the surface of the shock-tube wall. Therefore, under the initial pressure of a run, Ar gas laden with the desired proportion of N2H4 passed through the tube for at least 10 minutes, so that the gas in the tube was exchanged at least eight times. Checks with an interferometer (Haber-L6we, 1 m gas cell) had indicated that the interracial equilibrium was established at that point and that no appreciable conversion of N2H4 to NH3, N2, and H~ had taken place in the passage of hydrazine through the tube (molar refraction of gaseous N2H4 was determined to be R = 9.5 =t= 0.1 cm3/mole for white light; hence the accuracy of these determinations was about 2%). By this procedure, not only were errors in the N2H4 concentration circumvented, but also a certain separation of hydrazine from traces of water or from possible decomposition products, formed during storage of hydrazine in the saturator, could be achieved. NH3 is taken away with the first volumes of Ar passed through for flushing the tube (the same might hold for the possible decomposition product N2H~, which has a vapor pressure 14 similar to NH3). H20 is accumulated in the liquid phase. 15 Indeed, mass-spectrometric investigation of the vapor over commerical liquid hydrazine (Fluka AG, refractive index indicated ~ 2 % water) did not reveal any impurities which exceeded 0.1%. Thus, it appeared to be unnecessary to subject hydrazine to a further purification by distilling it under vacuum. When half of the hydrazine in the saturator vessels had evaporated, the remainder was discarded. Tank argon for welding, containing on an average 40 ppm 02 and 60 ppm N2, was used as a carrier gas. After further purification by leading it through a column filled with a catalyst (BTSKatalysator, furnished by BASF, Ludwigshafen), the O2 content could be decreased below the limit of detectability ( ~ 10 ppm). From this, the leak rate and the time elapsed between filling of the tube and the shock-wave experiment, the maximum oxygen content of the reaction gas was calculated to amount to ( 0 . 1 % with respect to
355
hydrazine, when the thermal decomposition was to be examined. Other gases, such as O2 and NH3, were obtained from commercial tanks without further purification.
Calculation of Shock-Wave Parameters The conditions behind the incident shocks were calculated from the instantaneous M a t h Number of the shock wave as it passed the observation windows. Because of the absence of cooling effects and boundary layers at this point, the conditions immediately behind the shock front are fixed by the simple laws of conservation. I t is only farther behind the shock that appreciable deviations from ideal shock-wave behavior occur, i.e., when the wake gases pile up in front of the observation window to form the reflected shock. An exact description of the cooling effect and of the deviation in reflected shock parameters by attenuation of the incident wave is extremely cumbersome. A rough estimate, however, indicates that the latter effects also cancel to a larger extent under the present conditions if the reflected-shock measurements are made at some distance from the closed end. This is actually borne out b y the consistency of data obtained in different tubes. (This cancellation does not hold any longer 16 if the reflected shock temperatures are above 3000~ The shock-wave parameters were obtained by an iteration procedure, l~ using tabulated ~7,1s enthalpy values which were appropriately combined with the enthalpies of reaction. Conditions behind the reflected shock can be computed exactly in a straightforward manner only when no reaction occurs or if the reaction time is very long as compared to the time the reflected shock needs to cover the distance between endplate and observation windows. During the course of the reaction, the flow conditions are not steady and the condition of quiescent gas behind the reflected shock does not apply. Depending upon the sign of the reaction enthalpy, the bulk of the reflected gases move away from or in the direction of the endplate. At the end of the reaction, however, the gases come to rest again. Any further motion would create expansion waves (or compression waves), the velocity of which equals the velocity of sound. Contrary to conditions in ChapmanJouguet detonations, these waves travel faster than the reflected-shock front, so that any flow of reacted gases would render the process unstable. If the reaction time is short as compared to the time which the reflected shock needs to cover the distance between endplate and observation
356
REACTION KINETICS
window (about 100 or 200 t~sec), the reflected shock becomes stationary again. Similar to a detonation, the equilibrated gas is preceded by a reaction zone and determines the reflected-shock velocity. Both conditions of the equilibrated gas and of the unreacted gas right behind the reflected-shock front in this case of "stationary reactive flow" are amenable to exact computation l~ again. The parameters of reflectedshock waves, the reaction time of which was of the same magnitude as the time needed for covering the distance between endplate and observation windows (unsteady case), was obtained by interpolation between the case of stationary reactive flow and the case without reaction. Ignoring these effects would lead to considerable errors even in the reflected-shock-front temperatures. For a mixture of 0 . 3 ~ N2tt4 with excess O2 in Ar, the shock-front temperature can be higher than calculated by as much as 80~ when not accounting for the coupling between reaction rate and flow properties. This method of calculation is confirmed by the observed tendency of reflected-shock speeds at the highest reaction rates of N2H4 to exceed theoretical values, which have been computed with the assumption of no reaction. The temperature difference across the reaction zone (30~ at slow reaction rates, 10~ at fast reaction rates for 0.3% N2H4 in Ar) was accounted for by a correction term to the initial reflectedshock-front temperature.
2500 ~_ they increase from about 7 X 10~ era2/ mole at 300~ to 2.5 X 105 cm2/mole at 1500~ Both increases occur in a non-Arrhenius fashion. For ammonia, the increase of absorption with temperature can be expressed b y e -- 0.029 exp (-- 2620 cm-1/kT) ~_2/mole at 2300~ and e ~ 0.020 exp (-- 6120 em-~/kT)A2/mole at
2500 ~.
Decomposition Rate of Hydrazine The recorded absorption traces of hydrazine (Fig. 1) indicate a complicated reaction. At the lower temperatures (1100~ < T5 < 1200~ and at the higher partial densities of N2H4 (p5 ~ 20 X 10-8 molelcm 3) noticeable decomposition does not start before a certain induction period has elapsed. The half-lives were most suitable for convenient characterization of the concentration-time curves in the whole range of initial partial densities and temperatures. These data are given in an Arrhenius type plot in Fig. 2 for experiments with approximately identical total densities of f)T ~ 7.5 X 10 -5 mole/cm a, the hydrazine concentration being varied by a factor of 10. At higher temperatures (T~ > 1300~ the decomposition seems to start out according to a first-order rate law and to pass over to a reaction phase with varying reaction order [-Fig. 3(a)],
i
I
I
I
I
I
,..~
Results
/
Extinction Coe~cients of Hydrazine and Ammonia For reliable evaluation of the absorption traces, and because they are interesting in themselves, a quantitative measurement of the absorption bands of hydrazine and its decomposition product ammonia had to be made. The results arc interpreted in terms of molar decadic extinction coefficients, ~. At higher temperatures the absorption bands of NH3 and N2H4 appear to be continuous. Within the accuracies of these measurements the extinction coefficients of both molecular species seem to be independent of geometric path length (x = 3.2 and 10 cm), of the nature of the carrier gas (He or At), of partial densities of the reactants, and of total pressure, such that straightforward evaluation of the absorption traces on the basis of Beer's law was facilitated. Whether measurements were made behind incident or reflected shocks was of no influence. For hydrazine, the extinction coefficients at 2 3 0 0 ~ increase from 1.8 X 105 cm2/mole at 300~ to 8.0 X 105 cm2/mole at 1500~ and at
't--7--)
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0.5
_
/ ~
#5--
/.0
_
=/'~/o
--
__/" I ~65
I
l
O,7O
O.75
I 080
t
0.85
I/T.10a(~ FIG. 2. Arrhenius plot of half-fives in N~H4 pyrolysis at total gas density of p~ = 7.5 X 10-5 mole/ 9 cm 3 (0.03%-0.5% N:H4 in At). -- -, p5 ~ 24 X 10-8 m o l e / c m 3 ; - - - , p5 ~ 8 X 10-8 mole/era3; . . . . . p5 2.5 X 10-8 mole/cm3; (p5 = partial density of reactant behind reflected shock prior to reaction).
357
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
polation of the log (log Io/I) vs t curves to the point where the reflected shock had crossed the recording lightbeam, initial absorption values of N2H4 could be determined up to 1600~ For experiments at 2300 .~ the light absorption of NH3, appearing as a reaction product, was accounted for in terms of a correction a which depended upon the fraction ~ of ammonia formed per mole N2I-I4 decomposed and amounted to < 1 0 % of the total absorption at the halflife period. Moreover, in the evaluation of all oscilloscope records with hydrazine, the initial absorption of N2H4 at room temperature had to be considered.
which order is lower, the smaller the initial hydrazine concentration (Fig. 3b). Oscilloscope records from runs involving temperatures above 1400~ were, in general, not evaluated according to the half-life method. Instead, for each individual absorption trace the compounded logarithm of the reciprocal of the light transmission Io/I (I0 = measured intensity of light without absorption from species in the shock tube) was plotted versus time. This allowed the evaluation of fast reactions even with halflives of 2 #sec, where the initial reaction phases were blurred out by schlieren effects, and limited response of the recording system. By linear extra-
reflected wove '=0
incident wove
incident
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.eflected wove t--O
wove
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2
4 6 8 10 12 14 16 reoction time [~SoC] (a)
0
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10 20 30. reoction time [j.secJ (b)
Fzo. 3(a). M = 2.46; ), = 2500 .~; 0.45% N2H4 in Ar; p5 = 26 X 10-s mole/cma; T6 = 1499~ kl' = 1.7 X 105 sec-1; sweep 20 psec/cm; sweep triggered twice so that incident shock was shown on the first and reflected shock on the second sweep. (b) M = 2.39; ), = 2300 .~; 0.034% N2H4 in Ar; p5 = 2.5 X 10-8 mole/cmS; Ts = 1415~ kl' = 0.76 X 10s sec-1; sweep 50 p s e c / c m .
358
REACTION KINETICS
55
""
5.0 b, ~\o
, ,~
IOg~ok
4.5!-
-
_
bx
the partial density of hydrazine (i.e., at low temperatures), they dwarf slight effects of the total gas density on the over-all reaction rate. At the higher temperatures, however, the first-order rate constants kt' vary slightly with the total gas density PT. This is illustrated by the Arrhenius plot in Fig. 4, which shows the temperature dependence of first-order rate constants abstracted from plots like Fig. 2 with the criteria: independence of partial N~H4 density and firstorder behavior from the start of the reaction (within the possible time resolution) to the halfperiod. One finds at
4~
7,5 X 10 -5 mole/em 3
PT =
k,' = 10,~-s exp E-(52.2 kcal/mole)/RT~
O.65
0.70
075
0.80
X sec-1 (1270~176
lIT. 103(OK)-~
FIG. 4. First-order rate constants in N2H4 pyrolysis at temperatures from 1260~ to 1600~ 0.030.5% N2H~ in Ar. The first-order rate constants kl' thus obtained, were expressed in terms of the halLlives; T89 = 0.693/k1' and included in Fig. 2 for comparison with results obtained at the lower temperatures or higher concentrations. Those points which are derived from oscilloscope records with first-order reaction behavior at least to the half-life are marked by filled symbols. Evidently, there is some sort of chain mechanism operating, and some consecutive steps lose in significance as one approaches higher temperatures and lower N2H4 concentrations. Above 1400~ M1 points obtained at the different partial densities of N2H4 merge together. Besides the generM dependence of half-lives on concentration, inspection of Fig. 3(a) indicates that the influence of some rate-determining consecutive step within the possible scheme of secondary reactions appears also in the log (log Io/I) vs t curves of each individual run. Deviation from a straight line illustrates the point where this consecutive step enables the whole chain mechanism to become operative. This event, naturally, takes place in a relatively later reaction stage the lower the partial density of N~H4, and depends on temperature. At 1400~ only the reaction with 2.5 X 10- s mole/cm 3 N~H4 showed first-order behavior during the decline in concentration b y a factor of 10 [Fig. 3(b)~.
The First-Order Rate Constant and Pressure Dependence When consecutive steps with chain character manifest themselves both b y an induction period and pronounced dependence of half-lives upon
and at
PT = 2.5 X 10-~ mole/em 3 h'
=
lo
exp r-(47.5
kcal/mole)/RT]
X sec-1 (1370~176 (Experimental uncertainties are probably larger than the evidently small standard deviation.) So far, these rate constants satisfy the criteria of a unimolecular reaction, even though they might incorporate the effect of some fast secondary reaction, the steady-state concentration of active species being established in such a short time that initial deviations from a first-order rate law are too small to be seen. One estimates that reliable evaluation of the concentrationtime curves was possible from the point at which the concentration had dropped b y 10%. If the unimolecular fission of N~H4
(1)
N2H4 --~ 2 NH2 0 ~ =2Z5"10"8; 0,30"/oNzH~ t
0/0/
(9 ?s =22.1"~~ ; 0.30"/,n.~H~.;om%0~ c~ Q ~ =2~9.~o-~ ; 029"/,N~,%;0,99*,'.0~ X / ~ (9 25 ~ r162=2~5.10.~ ; 0.2a'/.N~, ~oV'/.O~ . / ~ ~" 6+ ' _ / 0 ~..~
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oJ
25 --
r
/~" E=Z,3
1,5 --
~ __..,..".. - I .. ./.."
~Q."f 1.0 0.75
'
~ ~ 0.80
--
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I031T(OK) r r 0.85 0.90
Fie. 5. Arrhenius plot of half-lives in the oxidation of hydrazine behind reflected shock waves.
THERMAL
DECOMPOSITION
is followed by some fast consecutive steps, such as NH2 -{- N2H4 --* N2H3 + NHs,
(2)
it can be shown that the inequality k2 > 50(kJps) must be satisfied, in order that the influence of Reaction (2) upon the observed rate law escape detection by the described experimental technique. Taking some marginal data, one finds k2 > 7 X 1012 cm3/mole'sec. If this inequality should hold, the measured kl' would be an integral multiple of the real unimolecular rate constant kl.
Experiments with He as a Carrier Gas In order to test the inert-gas effect as well as to obtain some oscilloscope records without disturbing schlieren spikes, some experiments with He as a carrier gas were made. Measured half-lives ranged from 20 #sec (at 1370~ to 2.8 #sec (at 1540~ evaluation with some correction procedure for shock-front curvatureg). Variation of the partial density of N2H~ from 7 to 17 X 10-s mole/cm 3 was without material effect upon the observed first-order reaction rates. Total gas densities were about 3.4 X 10-6 mole/cm 3. The experimental scattering precluded the detection of finer inert-gas effects. Nonetheless, in general, the points coincided with the corresponding data obtained with Ar at the same conditions (agreement within 10%).
OF HYDRAZINE
359
pensated by inhibition of another one at these reaction temperatures. At higher oxygen concentrations the reaction obviously occurs according to a different mechanism than the pyrolysis. The absorption traces then start out with an induction period as in the pyrolysis, but fall off with a much more pronounced curvature. At temperatures higher than 1400~ however, even a tenfold excess of 02 over N2Ha does not result in an acceleration of the hydrazine decompsoition. This is also the point where the chain decomposition seems no longer capable of competing with the sequence of reactions which result in a first-order rate law. Hence, one might conclude that, under these conditions, 02 reacts not so much by primary attack onto hydrazine molecules (e.g., N2H4 + 02 ---> H02 + N2H3), but rather by making possible a fast secondary reaction. In view of the scarcity of data about interaction of O2 with radicals from the nitrogen-hydrogen system, the present results by themselves are insufficient to decide whether fast chain-propagating steps or chain-branching steps determine the oxidation of hydrazine. In any case, at these high temperatures and pressures, traces of O2 do not seem to falsify thermal-decomposition data.
Stoichiometry of the Thermal Decomposition of Hydrazine. For determining the amount of NH3
Oxidation of Hydrazine
formed in the pyrolysis of incident-reflected shock wave developed. This technique following processes, involving recording at 2300 ~:
N2Ht a combined technique has been is based on the spectrophotometric
The primary aim of these oxidation studies was to test whether traces of 02 occurring as impurities might affect the thermal decomposition rate of hydrazine. The high enthalpy of oxidation, AH ~ = --138.3 kcal/mole, involves appreciable deviations from an isothermal course of the reaction. Inasmuch as it is not known at what reaction stage the full heat of oxidation is liberated, the Arrhenius plot (Fig. 5) which exhibits the half-lives of mixtures with varying oxygen content may reflect only qualitatively the relationships of the oxidation reaction. Mixtures containing 0.3% N2H4 with oxygen concentrations of 0.22%, 0.99%, and 2.8% were examined. Data incorporated in Fig. 5 for the pyrolysis of pure hydrazine were obtained with mixture ratios N~I-I4:02 > 104. It is evident that small amounts of 02 do not accelerate the decomposition of hydrazine. To be sure, this does not exclude the possibility that oxygen might play an active part in the decomposition process. Sensitization of one decomposition step might be fortuitously corn-
N2H4 is decomposed behind incident shock waves (M > 3.0, T2 > ll00~ As the reflected shock travels through the wake of the incident shock after complete decomposition of N2H4, it raises the gas density by a factor of about 2 and the temperature to T~ > 2000~ At these temperatures, the extinction coefficient of one decomposition product, NH3, is high enough to produce an absorption signal at 2300 ~_, which can easily be evaluated 13 (Fig. 6). Since the extinction coefficient of NH3 had been determined in independent experiments up to temperatures of about 2900~ the amount of NH3 formed in the decomposition of N2H4 could be derived. At temperatures above 2100~ ammonia decomposes within the time of observation and the established rate law and rate constants verity unequivocally the identity of this species by comparison with other experiments (see below); (N2H2, which might have the same decomposition rate as NH3, has not been detected in significant quantities by DiesenS). The stoichiometry of the decomposition of
360
REACTION KINETICS
inc. refl.~ock ~)
........
~
J
.....
......
.- ...............
I: :"_'-"".
inc,
b)
um. =immw
....
Decomp. Decornp, of N;~H~, of NH3 c)
raft.shock
.......
Decornp. Decomp. of of NH~ d)
,:"
t
FIG. 6. N~H4 decomposition behind incident wave and NH3 decomposition behind reflected wave (f~ = degree of N2H4 conversion behind incident shock). (a) u~ = 0.962 mm//~sec; T2 = 1091~ T~ = 2150~ 0.09% N2H4; sweep 100 /~sec/cm. (b) u~ = 1.037 mm/gsec; T2 = 1214~ T5 = 2444~ 0.4% N~H4; T89 = 505 gsee; kFNH3-I = 2.6 X l0 s cm3/mole.sce; sweep 50 gsec/cm; ~ = 0.34; r = 0.80. (c) u~ = 1.047 mm/~tscc; T~ = 1235~ T5 = 2488~ 0.4% N2H4; T89 = 324 ~sec; k[-NH3~ = 5.0 X 10a cm3/mole-sec; sweep 50 gsec/cm; f~ = 0.65; v = 0.66. (d) ul = 1.094 mm/gsec; T~ = 1332~ T5 = 2730~ 0.12% N2H~; r89 = 119 gsee; kI-NHa] = 12.6 > l0 s cma/mole.sec; sweep 50 gsec/cm; ~ = 1.0; = 0.74. hydrazine was expressed in terms of the stoichiometric parameter v --- [-NH3-]/([-N2H4-]0 -[-N2H[]), viz., the ammonia formed divided b y the a m o u n t of hydrazine decomposed. The value of ~ ~ 1, established between l l 0 0 ~ and 1200~ was found to agree with results of other workers.3'7,s I n the larger tube, the distance between endplate and observation ~ n d o w was only 5 cm so that, at low reaction rates, the reflected shock was able to interact with the observed volume in wake of the incident wave before the reaction has gone to completion at this point (Fig. 6). , is found to be appreciably smaller, provided one
includes in [-N2H4-]o-[-N2H4~ only the hydrazine decomposed behind the incident wave, disregarding the residual material which is rapidly fragmented upon arrival of the reflected wave. Hence, only slowly decomposed N2H4 seems to form ammonia. If the temperature behind the incident shock exceeds 1400~ the reflected-shock temperature is close to 2900~ and ammonia decomposes too fast for safe determination of its initial concentration, so t h a t the described technique is no longer useful. Above 1400~ however, the extinction coefficient of NHa has become so high, that the quantities of NHs can be detected on the
refL shock inc.shock
ie~t i a l
sorption
---.
absorption behind inc. shock '~-- residual absorption due to NH3
FIG. 7. 1~2I:t4 decomposition behind reflected shock with residual absorption of NHa. ul = 0.809 mm/~sec; T5 = 1557~ 0.18% N2H G T89 ~ 2 ~sec; i, = 0.61; sweep 20 ttsec/cm I-triggered twice as in Fig. 3(a)].
361
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
tO
V ~5
II00 OK
J
I
1200~
1300~
I
I
1400~
1500 ~
I 1600~
Fio. 8. Production of NH3 in N2H4 pyrolysis. Decomposition behind incident wave
9 p~ = 11.6 0.32% N2H4 as = 5.0 0.40% N2H, ~)p~ = 2.2 0.085%N2H4 Op2 = 1.3 0.12%N2H4
pr pr pr pr
= = = =
0.7 1.2 2.5 1.2
Decomposition behind reflected wave
[ ] p5 = 11 0.44% N2H4 [] p5 = 11 0.18% H~H4 ---~p5 = 7.4 0.095%HsH4
pr = 2.5 pr = 5.7 p~, = 7.5
Molar partial densities of NsH4 behind incident wave (p,) and behind reflected wave (p~) in 10s mole/cm 8, total gas densities at reaction conditions of N:H4 (pr) in 105 mole/cm s. basis of their residual absorption after the complete decomposition of N2Ht. Then both processes, the decomposition of N2H4 and the determination of NHa can occur behind the reflected shock (Fig. 7). N:H4 decomposing at temperatures higher than 2000~ forms less than 10% NHs [-Fig. 6(a)7. In the region around 1400~ both methods furnish overlapping results. Figure 8 shows that there is no systematic dependence of p on the partial density of N2H4 or on total gas density. There is, however, a pronounced decline of v with temperature. If hydrazine is decomposed at 2000~ the quantity of NHs produced is so low that it cannot be detected any longer.
order character of the decomposition. The first-order rate constants were independent of partial density of NHs, which had been varied from 1.3 X 10-s to 150 X 10-8 mole/era s. They were, however, proportional to the total gas densities, varied between 1.2 and 10 X 10-5 mole/cm 3. Hence in Fig. 9, first-order rate constants divided b y the total gas densities (PT in mole/cm s) are plotted in Arrhenins fashion. This evidence strongly suggests that the unimolecular decomposition of NHs in the secondorder region of collisional activation has been observed. The rate law obeyed at temperatures between 2100 ~ and 2900~ and with Ar as a carrier gas reads:
The Pyrolysis of NH3. The experimental procedure and reduction of data were the same as with the NsI~ pyrolysis. The reaction mixtures were prepared by means of capillary flow meters, the one for NHa having been calibrated with the soap-bubble method. The wavelengths used varied between 2300 and 2500 ,~, depending upon the partial density of NH3 under investigation. Plotting of the compounded logarithm of the light transmission I o / I vs time, resulted in straight lines for all cases, establishing the first-
k = 1025.64 exp[.-- (79.5:l:2.5kcal/mole)/RT~ X cm3/mole" sec. This result is in rough agreement with data of other authors 2~ even though the evidence for the unimolecular character is much more cogent here, because of the wider range of conditions. The appearance of N H radicals during the decomposition of NHa at all temperatures above 2100~ suggests that these radicals play a major role in the mechanism.
362
REACTION
9.5
--
KINETICS
Reaction (1) as the primary decomposition mode of hydrazine and the first consecutive step (2) seem to be well established by these and other investigations 2.s:
*"
o o
,9.0
ee 9 .,r
NH2 + N2I-I4~ N~Ha + NHs.
(2)
.\
N~H~ -~ H + HNNH
\
~.5--
(1)
The chain mechanism, evidenced by the experimental results at lower temperatures, might include as the first propagating reactions:
~
tog k
N2H4 --~ 2 NH2
(3)
or
/ NH2 + N2H3 ~ NH3 + HNNH
NN
\ II
f
crn 3
7 9
7.5--d[Nt'r k fNH 7 [,4d I
0.35
I
(30
o~"\
NH~ + HNNH --. NHs + N2 + H.
,\kkk~ [
0.40 0.45 lIT.103(~ -'1
F m . 9. Arrhenius plot of second-order rate constants in NHs pyrolysis at temperature from 2100~
(5)
Combination of (3') and (5), excluding (3), would result in an ammonia production which is too high, even when further reasonable propagating or terminating reactions are taken into consideration, Spontaneous decomposition of N2H2 to N2 and H2 ['conceivable, if (3') forms the species
to 2900~ Discussion
In assigning a scheme of elementary reaction steps to the hydrazine decomposition under the described conditions, the following empirical facts have to be accounted for: (a) As one approaches higher temperatures and lower concentrations, the N2H4 decomposition becomes first-order in character. Elsewhere there seems to be an appreciable contribution of a chain mechanism, which gains in significance with increasing N2H4 content and gives rise to a pronounced induction period. (b) The ammonia produced at the lower temperatures corresponds closely to v = 1. With increasing temperature, v becomes smaller. At a reaction temperature of 2000~ no NHs could be detected with the described procedure and under the prevailing pressure conditions, v does not seem to depend very much upon the initial N2H4 concentration. A detailed discussion of possible elementary reactions is given elsewhere. 1~ Here, only the essential points are presented. The occurrence of
N N IH\ H
would not permit a reaction chain to develop. Hence (3) might be the dominant destruction reaction of N2H3, followed by H + N~H4 --* NHs + NH2
(4)
H + N2H4 ~ H2 -4- N2Ha
(4')
and the recombination reactions NH2 + NH~ -~ NH3 + NH
(6)
NH2 + NH2 --~ HNNH W H2
(6')
[M] + NH2 + H ~ NHs + [M].
(7)
If kt < k4f, practically no ammonia would be formed under conditions which favor chain decomposition (extrapolation of room temperature data 22,23to shock-wave conditions might not be quite iustified). By comparison with the H elimination from ethyl radicals, 24 the activation energy of (3) is estimated to be of the order of 42 keal/mole, v/z., (3) proceeds comparatively fast.
THERMAL DECOMPOSITION OF HYDRAZINE
Assuming that the measured k~' = kl and that chain terminating steps are not significant in the initial reaction phases, which are controlled only b y Reactions (1), (2), (3), and (4), an application of the steady-state approximations to the concentrations of N~H~ and H (by fitting the computed expression for the half-lives to experimental data) yields k~ ~ 1018'5exp [--- (17 kcal/mole)/RT~ X cm3/mole "sec. The theoretical curves derived with the above assumptions are shown in Fig. 2. An alternative interpretation on the basis of Reactions (1), (2), (3), and (4) might involve k2 ~ 10la cm3/mole-sec, ki' = 3kl and Reaction (3) being rate determining for the development of a chain mechanism. Numerical integration* of the four differential equations, however, has shown that this interpretation yields functions of half-lives, with respect to concentration and temperature, different from those observed experimentally. With kl' = kl and the above value for k2, scarcely any NH3 would be produced in chain propagating steps at temperatures above 1400~ In order to explain the experimental result v 0.5 at 1600~ with further decrease of v towards higher temperatures, the radical recombination (6'), followed by (5) is suggested. At temperatures above 2000~ Reaction (5) might be dwarfed by the unimolecular decomposition of H N N H HNNH --o NNH + H,
(8)
a reaction which m i g h t - - b y comparison with the NH3 decomposition--be expected to show strong dependence on total gas density, and which reduces the NH3 yields in the over-all decomposition. An appreciable contribution of Reaction (6) appears to be unlikely, because in that case the production of NH3 should be insensitive to variations in temperature. ACKNOWLEDGMENTS
We thank Professor W. Jost for his continuous interest and guidance in these investigations and the OAR for financial support through the European Office under Contract AF 61 (514)-1142. REFERENCES 1. AUDRIETH, L. F. AND OGG, B. A.: The Chemistry of ttydrazine: Wiley, 1951.
2. SZWARC, M.: Proc. Roy. Soc. (London) A198, 267 (1949). * The efficient cooperation of H. Eberius in solving this problem on an IBM 650 is gratefully acknowledged.
363
3. GRAY, P. AND LEE, I. C.: Seventh Symposium (International) on Combustion, p. 61, Butterworths, 1959. 4. HVSAIN, D. AND NORRISH, R. G. W.: Proc. Roy. Soc. (London) A273, 145 (1963). 5. GRAY, P., LEE, J. C., AND SPENCER, M.:
Combust. Flame 7, 320 (1963). 6. MOBERLu W. M.: J. Phys. Chem. 66, 366 (1962). 7. McHALE, E. T., KNOX, B. E., A~D PALMER, H. B.: Personal communication. 8. DIESEN, R. W.: J. Chem. Phys. 39, 2121 (1963). 9. MICHEL, K. W. AND WAGNER, I~. GG.: in Investigation of Gaseous Detonations and Shock-Wave Experiments with Hydrazine. Contract No. AF 61 (514)-1142, Technical Summary Report No. 2, 1960; MICHEL, K. W. AND WAGNER, H. GG. :
The Pyrolysis and Oxidation of Hydrazine in Shock Waves. Contract No. AF 61 (514)-1142, Technical Summary Report No. 3, 1962; MICHEL, K. W. AND WAGNER, H. GG.: in
Detonation and Shock-Tube Studies of Hydrazine and Nitrous Oxide. Contract No. AF 61 (514)1142, Annual Report No. 4, 1963. 10. MICHEL, K. W. AND WAGNER, H. GG.: Ber. der
Bunsen-Gesellschaft, in press. 11. MICHEL, K. W.: Ph.D. dissertation, G6ttingen, 1962. 12. RUDINGER, G.: Phys. Fluids $, 1463 (1961). 13. MICHEL, K. W. AND WAGNER, I-~. GG.: Z.
physik. Chemie NF 35, 392 (1962). 14. FONER, S. N. AND HUDSON, R. L.: J. Chem. Phys. 29, 442 (1958). 15. BURTLE, J. G.: Ind. Eng. Chem. ~4[, 1675 (1952). 16. MICHEL, K. W., OLSCHEWSKI, H. A., RICHTERINO, H., AND WAGNER, H. GG.: Z. physik. Chemie NF, in press. 17. Tables of Thermal Properties of Gases, National Bureau of Standards Circular 564, Washington, 1955. 18. LANDOLT--B6RNSTEIN: Kalorische ZustandsgrSssen, Vol. 2, Chap. 4, Springer, 1961. 19. BENNETT, m. G. AND DALBY, F. W.: J. Chem. Phys. 82, 1716 (1960). 20. MATHEWS, J. C., GIBBS, M. E., AND I~OLSEN,
J. N. : 139th National Meeting of the American Chemical Society, St. Louis, 1961. 21. JACOBS, T. A.: J. Phys. Chem. 67, 665 (1963). 22. BIRSE, E. A. B. AND MELVILLE, I-I. W.: Proc. Roy. Soc. (London) A175, 164 (1940). 23. SCHIAVELLO, M. AND VOLPI, G. G.: J. Chem. Phys. 37, 1510 (1962). 24. BrWATER, S. AND ROBERTS, R.: J. Chem. Phys. 19, 326 (1951). 25. DARWENT, D. DE B. AND ROBERTS, R.: Discussions Faraday Soc. 14, 55 (1953).
364
REACTION KINETICS
COMMENTS Prof. H. B. Palmer (Pennsylvania State University): Michel and I have had stimulating discussions about the compatibility of our studies, and it seems appropriate to mention what seem to me to be the critical points. The oscilloscope records of Michel and Wagner show induction periods t h a t are surely real. Their lengths correspond quite well to estimates of the times required to double the initial rate in the nonchain mechanism of our paper. I t is not obvious to me that, in their records, the rate increases by more than a factor of two during the induction period. If it really does, then our rejection of a chain must be wrong and we have to look for a source of error. The most obvious parameter to examine is the temperature in the reflected shock. There are two main questions a b o u t this:
temperatures above 1250~ Here, their experiments on the effect of the N2H4 concentration upon the half-life seem to show very beautifully the transition from a steady state for NH2 to a nonsteady state, i.e., from an over-all rate equal to twice the rate of the first step to a rate equal to one times it. A final point of interest on the mechanism is t h a t the nonchaln mechanism resembles t h a t for decomposition of hydrogen peroxide (Baldwin, 1%.,Jackson, D., Walker, R. W., and Webster, S. J. : this Symposinm p. 423). The parallels are: H202 --* 2 OH OH + H202 ~ H20 + HO2
N2H2 --* 2 NH2 NH~ + N2H~ --~ NH3 -t- N2H3
(a) How accurately known is the temperature of the initially-reflected shock? (b) How, if a t all, does the temperature in the reflected shock change with time?
The difference is:
In both shock-tube studies, ideal reflection is assumed. Although the shock tubes are somewhat different in geometry, errors from this assumption should be roughly the same in b o t h cases. Question (b) is, however, answered differently. Mainly by calculations, Michel concludes t h a t the temperature changes very little with time. From pressure records, we conclude t h a t it increases substantially during our dwell times. The pressure-increase phenomenon has been observed in at least four other laboratories and we are confident t h a t it is real. For example, it can be removed by modifying the flow in the shock tube [Tsang, W.: J. Chem. Phys. 40, 1498 (1964)]. When the pressure increase occurs, the gas must be heated in very nearly an isentropic manner. We do not find fault with Michel's calculations, b u t take rather the position t h a t it may not be correct to treat the reflected gas as a stagnant body adjacent to a cool wall (except at the endplate). We think t h a t aerodynamic effects are probably very significant in determining the temperature history. At short times, these temperature changes will be slight, which is why we have compared our data at low temperatures with Michel and Wagner's at
This makes us wonder whether due consideration has been given to the reaction OH q- HO2 -+ H~O + O2 as a possible third step in the decomposition of H202. I t gives the same kinetics as 2 HO2 ~ H202 + O.~, and should be very fast. Dr. K. W. Michel: I n Fig. 1, the rate of the decomposition of hydrazine increases b y more t h a n a factor of 2. I n general, absolute values of the initial reflected shock temperatures are not claimed to be more accurate, in this case, t h a n to about 30~ even though the consistency of relative temperatures which is material for the accuracy of the activation energy) is much better. This is also borne out by the excellent agreement of data obtained in different shock tubes of this investigation. The temperature increase of the reflected gas, due to adiabatic postcompression, gains in significance at reaction times in excess of 500 usec. Again, this is substantiated experimentally b y the observation of first-order rate laws over tong reaction periods in the decompositions of NH~, COs (Ref. 16) and N20 [JosT, W., MICHEL, K. W., TROE, J., AND WAGNER, H. Go.: Z. Naturforschung 19a, 59 (1964)] with the same equipment.
2 HO2 -o H202 + O~
N2H~ ~ NH~ --* NH~ -t- (NzH2) --+ N2 ~ H~.
pp. 353-364,
The Combustion'Institute, 1965-
THE PYROLYSIS AND OXIDATION OF H Y D R A Z I N E B E H I N D S H O C K WAVES K. W. MICHEL, AND H. GG. WAGNER
Institut fiir physikalische Chemie der Universit~t GOltingen, G6ttingen, Germany The thermal decomposition of hydrazinc highly diluted with Ar or He (0.03%-0.5% N2H4) has been investigated spectrophotometricaUy behind shock waves in the temperature range of 1100~ to 1600~ At the lower temperatures and higher concentrations of N2It~, the absorption traces recordcd on the oscilloscope exhibit an induction period and a certain dependence of the half-lives of N~H4 on concentration, viz., the characteristics of a chain decomposition. As one approaches higher temperatures (T > 1300~ and lower partial densities of N_~H4, the rate law for the decomposition becomes roughly first order; the first-order rate constant being independent of N2H4 concentration but revealing a slight dependence on total gas density pv: log k ~ 12.8 - [-(52 kcal/mole)/2.3RT] at or = 7.5 X 10-s inole/cm~ and log k ~ 12.0 - [(48 kcal/mole)/2.3RT] at pT = 2.5 X 10-6 mole/era 3 (k in see-0. The addition of oxygen to an 0.3% N2H~-Ar mixture causes noticeablc acceleration of the decomposition only when added in appreciable amounts ([-O2-]/[N~H,-] > 2). The amount of NHs formed during the pyrolysis of N2H4 has been de~rmined by a combined incident-reflected shock-wave technique, again using uv absorption at 2300 A for quantitative determination of both reactant and product. The ratio rNHz-](formed)/[-N;H~-](decomposed) decreases from 1 at 1100~ to 0.5 at 1600~ and is about 0 at 2000~ At temperatures above 1400~ the NH radical could be detected photometrically. Introduction The investigation of the thermal decomposition of gaseous hydrazine by classical methods suffers greatly from wall catalysis. 1,2 To obviate this difficulty, other techniques have been resorted to such as decomposition flames, 8 flash photolysis, 4 studies of explosion characteristics. 5 Information about the controlling processes under nearly isothermal conditions, however, may be obtained most conveniently by means of shockwave techniques. The single-pulse method has been applied to the N2H4 pyrolysis by Moberly e and by McHale et al. 7 in the temperature range 930~176 and 950~176 respectively. The data of the latter investigation has been interpreted in terms of a first-order rate law, giving k = 1.1 X 10~3 exp ( - 5 5 kcal/RT)sec -1. Mass spectroscopy has been combined with the shock-wave technique by Dicsen s in a study of the N2H4 pyrolysis and at temperatures from 1200~ to 2500~ at appreciably lower pressures
(total gas density, pr = 7 X 10-7 to 13 X 10-~ mole/era 3) than in the above-mentioned investigations and in this one (PT = 2.5 X 10- ~ to 14 X 10-5 mole/cm3). In this study a uv-absorption technique was used, permitting the continuous recording of the kinetic behavior of hydrazinc, which was present in comparatively low concentrations (less than 0.5% N2H4 in Ar or He). Supplementary information could be derived by tracing the concentration course of the NIt radical and by determining the amount of NHs produced in tile N~I/4 pyrolysis. The primary goal in the oxidation experiments was to test the effect of small amounts of oxygen upon the pyrolysis. Apparatus and Method Three different shock tubes, each made of aluminum, wcrc used, in order to eliminate the effects of special properties of the apparatus (boundary-layer effects, attenuation, catalytic
353
354
REACTION
KINETICS
Fro. 1. N2H4 pyrolysis observed at 2300 .~. M = 2.11; re~u1 = 0.596; ,Au/u~ = 3.0%/m; 0.28% N2H4 in Ar; sweep 200 t~see/cm; p5 = 24 )< 10s mole/cm3; pT = 8.6 )< 10-s mole/cmS; Ts = l l20~ decomposition of hydrazine prior to shock-wave experiment, etc.) upon the results. Two square shock tubes, 3.2 cm i.d., one being used without any further treatment and the other having been eloxized, each consisting of a 180-cm-long highpressure section and a 250-cm-long test section. Thus, reaction times behind reflected shocks of at least 1 msec were possible (longer ones were not considered suitable because of deviations from ideal shock behavior such as cooling from the walls and boundary layer effects) ~e. Ten cm from the reflecting endplate, quartz windows (Ultrasil, 90% transmission at 2300 ~,) were inserted flush with the walls for spectrophotometric recording with uv light. The third tube with a round cross section of 10 cm i.d. had a 280-cm-long high-pressure section and a 420-cm-long test section. Quartz windows were inserted 5 cm from the closed end. The leak rate of the low-pressure section was not more than 10--5 mm IIg/min. The tubes were equipped with separate intake and pumping valves, so that the reaction mixture, prepared in a glass system with Teflon joints, could be flushed through the tube. for a length of time before each experiment. A schlieren set-up was used for measuring the speed and attenuation of the incident shock waves, such that the instantaneous shock velocity was measured to within 0.2%. The shock-wave attenuation in the large tube varied between 0 and 1.6%/m, in the small one it was about 3.6%/m for both carrier gases, Ar and tie. Reflected shock speeds could be determined to 1.5% accuracy and provided, together with the attenuation values and the correspondence oi the measured Mach Number with the initial pressure ratio of driver and test gas (within 30% of theoretical), a convenient control for the stationary character of the shock waves. An effective M a t h Number had been defined to account for deviations from ideal shock
r] = 445 gsec.
behavior and proved to be consistent when comparing results from different tubes and from shocks with different attenuation. An almost parallel beam of ultraviolet light from a Xe high-pressure arc (Osram, XBO 150 W) passed through the quartz windows near the end of the test scctioa and onto the slit of a monochromator (Zciss M4 Q II; an EMI-6256 A photomultiplier was used as detector). The incident-light intensity was measured before and after each run with a rotating scctor of 2.8 kcps. The electrical time constant of the recording lmit amounted to 0.4 or 1 ttsec. Measurements of the N2tI4 and NHa concentrations were made at 2300 and 2500 ~, chosen so that the maximum absorption always fell into the most favorable range of 10% to 70%. The absorption traces were evaluated with an accuracy of 5%. Figure 1 shows a typical oscilloscope record of the absorption of hydrazine. The first step in the absorption trace occurred with the passage of the incident-shock front. The trace then remained constant for some 400 t~sec, indicating that there was no appreciable decomposition of hydrazine within 1 msec (particle time scale) at 700~ Even behind the reflected shock, fast decomposition of hydrazine did not occur before a certain induction period had elapsed. The data for characterizing an experiment are given in the legend: M is the effective Mach Number of the incident shock, A u / u its attenuation per m, pT the total gas density at reaction conditions, T5 the calculated temperature behind the reflected shock, and r89 the measured half-life of hydrazine. Further experimental details have been described elsewhere. ~n'13
Reactants For preparing N2H4-Ar mixtures, part of the Ar-stream was bubbled (at known pressure)
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
through liquid hydrazine in a double saturator at 32~ Because N2H4 evaporates sluggishly, this was not enough for accurately establishing the concentration, and the gas laden with N2H4 vapor passed through a condenser, the temperature of which was kept constant at 22.0 ~ -40.05~ Beyond the condenser, the Ar stream saturated with hydrazine was reunited with the main stream of pure Ar or Ar-02 mixtures. The ratio of flow velocities was fixed by a set of properly gauged capillaries or by Bodenstein valves. N2H4 is strongly adsorbed by the acidic aluminum oxide at the surface of the shock-tube wall. Therefore, under the initial pressure of a run, Ar gas laden with the desired proportion of N2H4 passed through the tube for at least 10 minutes, so that the gas in the tube was exchanged at least eight times. Checks with an interferometer (Haber-L6we, 1 m gas cell) had indicated that the interracial equilibrium was established at that point and that no appreciable conversion of N2H4 to NH3, N2, and H~ had taken place in the passage of hydrazine through the tube (molar refraction of gaseous N2H4 was determined to be R = 9.5 =t= 0.1 cm3/mole for white light; hence the accuracy of these determinations was about 2%). By this procedure, not only were errors in the N2H4 concentration circumvented, but also a certain separation of hydrazine from traces of water or from possible decomposition products, formed during storage of hydrazine in the saturator, could be achieved. NH3 is taken away with the first volumes of Ar passed through for flushing the tube (the same might hold for the possible decomposition product N2H~, which has a vapor pressure 14 similar to NH3). H20 is accumulated in the liquid phase. 15 Indeed, mass-spectrometric investigation of the vapor over commerical liquid hydrazine (Fluka AG, refractive index indicated ~ 2 % water) did not reveal any impurities which exceeded 0.1%. Thus, it appeared to be unnecessary to subject hydrazine to a further purification by distilling it under vacuum. When half of the hydrazine in the saturator vessels had evaporated, the remainder was discarded. Tank argon for welding, containing on an average 40 ppm 02 and 60 ppm N2, was used as a carrier gas. After further purification by leading it through a column filled with a catalyst (BTSKatalysator, furnished by BASF, Ludwigshafen), the O2 content could be decreased below the limit of detectability ( ~ 10 ppm). From this, the leak rate and the time elapsed between filling of the tube and the shock-wave experiment, the maximum oxygen content of the reaction gas was calculated to amount to ( 0 . 1 % with respect to
355
hydrazine, when the thermal decomposition was to be examined. Other gases, such as O2 and NH3, were obtained from commercial tanks without further purification.
Calculation of Shock-Wave Parameters The conditions behind the incident shocks were calculated from the instantaneous M a t h Number of the shock wave as it passed the observation windows. Because of the absence of cooling effects and boundary layers at this point, the conditions immediately behind the shock front are fixed by the simple laws of conservation. I t is only farther behind the shock that appreciable deviations from ideal shock-wave behavior occur, i.e., when the wake gases pile up in front of the observation window to form the reflected shock. An exact description of the cooling effect and of the deviation in reflected shock parameters by attenuation of the incident wave is extremely cumbersome. A rough estimate, however, indicates that the latter effects also cancel to a larger extent under the present conditions if the reflected-shock measurements are made at some distance from the closed end. This is actually borne out b y the consistency of data obtained in different tubes. (This cancellation does not hold any longer 16 if the reflected shock temperatures are above 3000~ The shock-wave parameters were obtained by an iteration procedure, l~ using tabulated ~7,1s enthalpy values which were appropriately combined with the enthalpies of reaction. Conditions behind the reflected shock can be computed exactly in a straightforward manner only when no reaction occurs or if the reaction time is very long as compared to the time the reflected shock needs to cover the distance between endplate and observation windows. During the course of the reaction, the flow conditions are not steady and the condition of quiescent gas behind the reflected shock does not apply. Depending upon the sign of the reaction enthalpy, the bulk of the reflected gases move away from or in the direction of the endplate. At the end of the reaction, however, the gases come to rest again. Any further motion would create expansion waves (or compression waves), the velocity of which equals the velocity of sound. Contrary to conditions in ChapmanJouguet detonations, these waves travel faster than the reflected-shock front, so that any flow of reacted gases would render the process unstable. If the reaction time is short as compared to the time which the reflected shock needs to cover the distance between endplate and observation
356
REACTION KINETICS
window (about 100 or 200 t~sec), the reflected shock becomes stationary again. Similar to a detonation, the equilibrated gas is preceded by a reaction zone and determines the reflected-shock velocity. Both conditions of the equilibrated gas and of the unreacted gas right behind the reflected-shock front in this case of "stationary reactive flow" are amenable to exact computation l~ again. The parameters of reflectedshock waves, the reaction time of which was of the same magnitude as the time needed for covering the distance between endplate and observation windows (unsteady case), was obtained by interpolation between the case of stationary reactive flow and the case without reaction. Ignoring these effects would lead to considerable errors even in the reflected-shock-front temperatures. For a mixture of 0 . 3 ~ N2tt4 with excess O2 in Ar, the shock-front temperature can be higher than calculated by as much as 80~ when not accounting for the coupling between reaction rate and flow properties. This method of calculation is confirmed by the observed tendency of reflected-shock speeds at the highest reaction rates of N2H4 to exceed theoretical values, which have been computed with the assumption of no reaction. The temperature difference across the reaction zone (30~ at slow reaction rates, 10~ at fast reaction rates for 0.3% N2H4 in Ar) was accounted for by a correction term to the initial reflectedshock-front temperature.
2500 ~_ they increase from about 7 X 10~ era2/ mole at 300~ to 2.5 X 105 cm2/mole at 1500~ Both increases occur in a non-Arrhenius fashion. For ammonia, the increase of absorption with temperature can be expressed b y e -- 0.029 exp (-- 2620 cm-1/kT) ~_2/mole at 2300~ and e ~ 0.020 exp (-- 6120 em-~/kT)A2/mole at
2500 ~.
Decomposition Rate of Hydrazine The recorded absorption traces of hydrazine (Fig. 1) indicate a complicated reaction. At the lower temperatures (1100~ < T5 < 1200~ and at the higher partial densities of N2H4 (p5 ~ 20 X 10-8 molelcm 3) noticeable decomposition does not start before a certain induction period has elapsed. The half-lives were most suitable for convenient characterization of the concentration-time curves in the whole range of initial partial densities and temperatures. These data are given in an Arrhenius type plot in Fig. 2 for experiments with approximately identical total densities of f)T ~ 7.5 X 10 -5 mole/cm a, the hydrazine concentration being varied by a factor of 10. At higher temperatures (T~ > 1300~ the decomposition seems to start out according to a first-order rate law and to pass over to a reaction phase with varying reaction order [-Fig. 3(a)],
i
I
I
I
I
I
,..~
Results
/
Extinction Coe~cients of Hydrazine and Ammonia For reliable evaluation of the absorption traces, and because they are interesting in themselves, a quantitative measurement of the absorption bands of hydrazine and its decomposition product ammonia had to be made. The results arc interpreted in terms of molar decadic extinction coefficients, ~. At higher temperatures the absorption bands of NH3 and N2H4 appear to be continuous. Within the accuracies of these measurements the extinction coefficients of both molecular species seem to be independent of geometric path length (x = 3.2 and 10 cm), of the nature of the carrier gas (He or At), of partial densities of the reactants, and of total pressure, such that straightforward evaluation of the absorption traces on the basis of Beer's law was facilitated. Whether measurements were made behind incident or reflected shocks was of no influence. For hydrazine, the extinction coefficients at 2 3 0 0 ~ increase from 1.8 X 105 cm2/mole at 300~ to 8.0 X 105 cm2/mole at 1500~ and at
't--7--)
'"
,;,~."'o
W4;0o o
0.5
_
/ ~
#5--
/.0
_
=/'~/o
--
__/" I ~65
I
l
O,7O
O.75
I 080
t
0.85
I/T.10a(~ FIG. 2. Arrhenius plot of half-fives in N~H4 pyrolysis at total gas density of p~ = 7.5 X 10-5 mole/ 9 cm 3 (0.03%-0.5% N:H4 in At). -- -, p5 ~ 24 X 10-8 m o l e / c m 3 ; - - - , p5 ~ 8 X 10-8 mole/era3; . . . . . p5 2.5 X 10-8 mole/cm3; (p5 = partial density of reactant behind reflected shock prior to reaction).
357
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
polation of the log (log Io/I) vs t curves to the point where the reflected shock had crossed the recording lightbeam, initial absorption values of N2H4 could be determined up to 1600~ For experiments at 2300 .~ the light absorption of NH3, appearing as a reaction product, was accounted for in terms of a correction a which depended upon the fraction ~ of ammonia formed per mole N2I-I4 decomposed and amounted to < 1 0 % of the total absorption at the halflife period. Moreover, in the evaluation of all oscilloscope records with hydrazine, the initial absorption of N2H4 at room temperature had to be considered.
which order is lower, the smaller the initial hydrazine concentration (Fig. 3b). Oscilloscope records from runs involving temperatures above 1400~ were, in general, not evaluated according to the half-life method. Instead, for each individual absorption trace the compounded logarithm of the reciprocal of the light transmission Io/I (I0 = measured intensity of light without absorption from species in the shock tube) was plotted versus time. This allowed the evaluation of fast reactions even with halflives of 2 #sec, where the initial reaction phases were blurred out by schlieren effects, and limited response of the recording system. By linear extra-
reflected wove '=0
incident wove
incident
\
:-.'
.eflected wove t--O
wove
"t
i I I I - 1.6
off l/
1.4
I
I
|11
L
.
*
,
i
*
*
*
,
IT I i i 2+
IOg(Iog ~ / I " o: )
7.2
*..
1.0 O.8 O.6 O.4
O2 0
0
2
4 6 8 10 12 14 16 reoction time [~SoC] (a)
0
I
I
I
10 20 30. reoction time [j.secJ (b)
Fzo. 3(a). M = 2.46; ), = 2500 .~; 0.45% N2H4 in Ar; p5 = 26 X 10-s mole/cma; T6 = 1499~ kl' = 1.7 X 105 sec-1; sweep 20 psec/cm; sweep triggered twice so that incident shock was shown on the first and reflected shock on the second sweep. (b) M = 2.39; ), = 2300 .~; 0.034% N2H4 in Ar; p5 = 2.5 X 10-8 mole/cmS; Ts = 1415~ kl' = 0.76 X 10s sec-1; sweep 50 p s e c / c m .
358
REACTION KINETICS
55
""
5.0 b, ~\o
, ,~
IOg~ok
4.5!-
-
_
bx
the partial density of hydrazine (i.e., at low temperatures), they dwarf slight effects of the total gas density on the over-all reaction rate. At the higher temperatures, however, the first-order rate constants kt' vary slightly with the total gas density PT. This is illustrated by the Arrhenius plot in Fig. 4, which shows the temperature dependence of first-order rate constants abstracted from plots like Fig. 2 with the criteria: independence of partial N~H4 density and firstorder behavior from the start of the reaction (within the possible time resolution) to the halfperiod. One finds at
4~
7,5 X 10 -5 mole/em 3
PT =
k,' = 10,~-s exp E-(52.2 kcal/mole)/RT~
O.65
0.70
075
0.80
X sec-1 (1270~176
lIT. 103(OK)-~
FIG. 4. First-order rate constants in N2H4 pyrolysis at temperatures from 1260~ to 1600~ 0.030.5% N2H~ in Ar. The first-order rate constants kl' thus obtained, were expressed in terms of the halLlives; T89 = 0.693/k1' and included in Fig. 2 for comparison with results obtained at the lower temperatures or higher concentrations. Those points which are derived from oscilloscope records with first-order reaction behavior at least to the half-life are marked by filled symbols. Evidently, there is some sort of chain mechanism operating, and some consecutive steps lose in significance as one approaches higher temperatures and lower N2H4 concentrations. Above 1400~ M1 points obtained at the different partial densities of N2H4 merge together. Besides the generM dependence of half-lives on concentration, inspection of Fig. 3(a) indicates that the influence of some rate-determining consecutive step within the possible scheme of secondary reactions appears also in the log (log Io/I) vs t curves of each individual run. Deviation from a straight line illustrates the point where this consecutive step enables the whole chain mechanism to become operative. This event, naturally, takes place in a relatively later reaction stage the lower the partial density of N~H4, and depends on temperature. At 1400~ only the reaction with 2.5 X 10- s mole/cm 3 N~H4 showed first-order behavior during the decline in concentration b y a factor of 10 [Fig. 3(b)~.
The First-Order Rate Constant and Pressure Dependence When consecutive steps with chain character manifest themselves both b y an induction period and pronounced dependence of half-lives upon
and at
PT = 2.5 X 10-~ mole/em 3 h'
=
lo
exp r-(47.5
kcal/mole)/RT]
X sec-1 (1370~176 (Experimental uncertainties are probably larger than the evidently small standard deviation.) So far, these rate constants satisfy the criteria of a unimolecular reaction, even though they might incorporate the effect of some fast secondary reaction, the steady-state concentration of active species being established in such a short time that initial deviations from a first-order rate law are too small to be seen. One estimates that reliable evaluation of the concentrationtime curves was possible from the point at which the concentration had dropped b y 10%. If the unimolecular fission of N~H4
(1)
N2H4 --~ 2 NH2 0 ~ =2Z5"10"8; 0,30"/oNzH~ t
0/0/
(9 ?s =22.1"~~ ; 0.30"/,n.~H~.;om%0~ c~ Q ~ =2~9.~o-~ ; 029"/,N~,%;0,99*,'.0~ X / ~ (9 25 ~ r162=2~5.10.~ ; 0.2a'/.N~, ~oV'/.O~ . / ~ ~" 6+ ' _ / 0 ~..~
,o,.
oJ
25 --
r
/~" E=Z,3
1,5 --
~ __..,..".. - I .. ./.."
~Q."f 1.0 0.75
'
~ ~ 0.80
--
"
I031T(OK) r r 0.85 0.90
Fie. 5. Arrhenius plot of half-lives in the oxidation of hydrazine behind reflected shock waves.
THERMAL
DECOMPOSITION
is followed by some fast consecutive steps, such as NH2 -{- N2H4 --* N2H3 + NHs,
(2)
it can be shown that the inequality k2 > 50(kJps) must be satisfied, in order that the influence of Reaction (2) upon the observed rate law escape detection by the described experimental technique. Taking some marginal data, one finds k2 > 7 X 1012 cm3/mole'sec. If this inequality should hold, the measured kl' would be an integral multiple of the real unimolecular rate constant kl.
Experiments with He as a Carrier Gas In order to test the inert-gas effect as well as to obtain some oscilloscope records without disturbing schlieren spikes, some experiments with He as a carrier gas were made. Measured half-lives ranged from 20 #sec (at 1370~ to 2.8 #sec (at 1540~ evaluation with some correction procedure for shock-front curvatureg). Variation of the partial density of N2H~ from 7 to 17 X 10-s mole/cm 3 was without material effect upon the observed first-order reaction rates. Total gas densities were about 3.4 X 10-6 mole/cm 3. The experimental scattering precluded the detection of finer inert-gas effects. Nonetheless, in general, the points coincided with the corresponding data obtained with Ar at the same conditions (agreement within 10%).
OF HYDRAZINE
359
pensated by inhibition of another one at these reaction temperatures. At higher oxygen concentrations the reaction obviously occurs according to a different mechanism than the pyrolysis. The absorption traces then start out with an induction period as in the pyrolysis, but fall off with a much more pronounced curvature. At temperatures higher than 1400~ however, even a tenfold excess of 02 over N2Ha does not result in an acceleration of the hydrazine decompsoition. This is also the point where the chain decomposition seems no longer capable of competing with the sequence of reactions which result in a first-order rate law. Hence, one might conclude that, under these conditions, 02 reacts not so much by primary attack onto hydrazine molecules (e.g., N2H4 + 02 ---> H02 + N2H3), but rather by making possible a fast secondary reaction. In view of the scarcity of data about interaction of O2 with radicals from the nitrogen-hydrogen system, the present results by themselves are insufficient to decide whether fast chain-propagating steps or chain-branching steps determine the oxidation of hydrazine. In any case, at these high temperatures and pressures, traces of O2 do not seem to falsify thermal-decomposition data.
Stoichiometry of the Thermal Decomposition of Hydrazine. For determining the amount of NH3
Oxidation of Hydrazine
formed in the pyrolysis of incident-reflected shock wave developed. This technique following processes, involving recording at 2300 ~:
N2Ht a combined technique has been is based on the spectrophotometric
The primary aim of these oxidation studies was to test whether traces of 02 occurring as impurities might affect the thermal decomposition rate of hydrazine. The high enthalpy of oxidation, AH ~ = --138.3 kcal/mole, involves appreciable deviations from an isothermal course of the reaction. Inasmuch as it is not known at what reaction stage the full heat of oxidation is liberated, the Arrhenius plot (Fig. 5) which exhibits the half-lives of mixtures with varying oxygen content may reflect only qualitatively the relationships of the oxidation reaction. Mixtures containing 0.3% N2H4 with oxygen concentrations of 0.22%, 0.99%, and 2.8% were examined. Data incorporated in Fig. 5 for the pyrolysis of pure hydrazine were obtained with mixture ratios N~I-I4:02 > 104. It is evident that small amounts of 02 do not accelerate the decomposition of hydrazine. To be sure, this does not exclude the possibility that oxygen might play an active part in the decomposition process. Sensitization of one decomposition step might be fortuitously corn-
N2H4 is decomposed behind incident shock waves (M > 3.0, T2 > ll00~ As the reflected shock travels through the wake of the incident shock after complete decomposition of N2H4, it raises the gas density by a factor of about 2 and the temperature to T~ > 2000~ At these temperatures, the extinction coefficient of one decomposition product, NH3, is high enough to produce an absorption signal at 2300 ~_, which can easily be evaluated 13 (Fig. 6). Since the extinction coefficient of NH3 had been determined in independent experiments up to temperatures of about 2900~ the amount of NH3 formed in the decomposition of N2H4 could be derived. At temperatures above 2100~ ammonia decomposes within the time of observation and the established rate law and rate constants verity unequivocally the identity of this species by comparison with other experiments (see below); (N2H2, which might have the same decomposition rate as NH3, has not been detected in significant quantities by DiesenS). The stoichiometry of the decomposition of
360
REACTION KINETICS
inc. refl.~ock ~)
........
~
J
.....
......
.- ...............
I: :"_'-"".
inc,
b)
um. =immw
....
Decomp. Decornp, of N;~H~, of NH3 c)
raft.shock
.......
Decornp. Decomp. of of NH~ d)
,:"
t
FIG. 6. N~H4 decomposition behind incident wave and NH3 decomposition behind reflected wave (f~ = degree of N2H4 conversion behind incident shock). (a) u~ = 0.962 mm//~sec; T2 = 1091~ T~ = 2150~ 0.09% N2H4; sweep 100 /~sec/cm. (b) u~ = 1.037 mm/gsec; T2 = 1214~ T5 = 2444~ 0.4% N~H4; T89 = 505 gsee; kFNH3-I = 2.6 X l0 s cm3/mole.sce; sweep 50 gsec/cm; ~ = 0.34; r = 0.80. (c) u~ = 1.047 mm/~tscc; T~ = 1235~ T5 = 2488~ 0.4% N2H4; T89 = 324 ~sec; k[-NH3~ = 5.0 X 10a cm3/mole-sec; sweep 50 gsec/cm; f~ = 0.65; v = 0.66. (d) ul = 1.094 mm/gsec; T~ = 1332~ T5 = 2730~ 0.12% N2H~; r89 = 119 gsee; kI-NHa] = 12.6 > l0 s cma/mole.sec; sweep 50 gsec/cm; ~ = 1.0; = 0.74. hydrazine was expressed in terms of the stoichiometric parameter v --- [-NH3-]/([-N2H4-]0 -[-N2H[]), viz., the ammonia formed divided b y the a m o u n t of hydrazine decomposed. The value of ~ ~ 1, established between l l 0 0 ~ and 1200~ was found to agree with results of other workers.3'7,s I n the larger tube, the distance between endplate and observation ~ n d o w was only 5 cm so that, at low reaction rates, the reflected shock was able to interact with the observed volume in wake of the incident wave before the reaction has gone to completion at this point (Fig. 6). , is found to be appreciably smaller, provided one
includes in [-N2H4-]o-[-N2H4~ only the hydrazine decomposed behind the incident wave, disregarding the residual material which is rapidly fragmented upon arrival of the reflected wave. Hence, only slowly decomposed N2H4 seems to form ammonia. If the temperature behind the incident shock exceeds 1400~ the reflected-shock temperature is close to 2900~ and ammonia decomposes too fast for safe determination of its initial concentration, so t h a t the described technique is no longer useful. Above 1400~ however, the extinction coefficient of NHa has become so high, that the quantities of NHs can be detected on the
refL shock inc.shock
ie~t i a l
sorption
---.
absorption behind inc. shock '~-- residual absorption due to NH3
FIG. 7. 1~2I:t4 decomposition behind reflected shock with residual absorption of NHa. ul = 0.809 mm/~sec; T5 = 1557~ 0.18% N2H G T89 ~ 2 ~sec; i, = 0.61; sweep 20 ttsec/cm I-triggered twice as in Fig. 3(a)].
361
T H E R M A L D E C O M P O S I T I O N OF H Y D R A Z I N E
tO
V ~5
II00 OK
J
I
1200~
1300~
I
I
1400~
1500 ~
I 1600~
Fio. 8. Production of NH3 in N2H4 pyrolysis. Decomposition behind incident wave
9 p~ = 11.6 0.32% N2H4 as = 5.0 0.40% N2H, ~)p~ = 2.2 0.085%N2H4 Op2 = 1.3 0.12%N2H4
pr pr pr pr
= = = =
0.7 1.2 2.5 1.2
Decomposition behind reflected wave
[ ] p5 = 11 0.44% N2H4 [] p5 = 11 0.18% H~H4 ---~p5 = 7.4 0.095%HsH4
pr = 2.5 pr = 5.7 p~, = 7.5
Molar partial densities of NsH4 behind incident wave (p,) and behind reflected wave (p~) in 10s mole/cm 8, total gas densities at reaction conditions of N:H4 (pr) in 105 mole/cm s. basis of their residual absorption after the complete decomposition of N2Ht. Then both processes, the decomposition of N2H4 and the determination of NHa can occur behind the reflected shock (Fig. 7). N:H4 decomposing at temperatures higher than 2000~ forms less than 10% NHs [-Fig. 6(a)7. In the region around 1400~ both methods furnish overlapping results. Figure 8 shows that there is no systematic dependence of p on the partial density of N2H4 or on total gas density. There is, however, a pronounced decline of v with temperature. If hydrazine is decomposed at 2000~ the quantity of NHs produced is so low that it cannot be detected any longer.
order character of the decomposition. The first-order rate constants were independent of partial density of NHs, which had been varied from 1.3 X 10-s to 150 X 10-8 mole/era s. They were, however, proportional to the total gas densities, varied between 1.2 and 10 X 10-5 mole/cm 3. Hence in Fig. 9, first-order rate constants divided b y the total gas densities (PT in mole/cm s) are plotted in Arrhenins fashion. This evidence strongly suggests that the unimolecular decomposition of NHs in the secondorder region of collisional activation has been observed. The rate law obeyed at temperatures between 2100 ~ and 2900~ and with Ar as a carrier gas reads:
The Pyrolysis of NH3. The experimental procedure and reduction of data were the same as with the NsI~ pyrolysis. The reaction mixtures were prepared by means of capillary flow meters, the one for NHa having been calibrated with the soap-bubble method. The wavelengths used varied between 2300 and 2500 ,~, depending upon the partial density of NH3 under investigation. Plotting of the compounded logarithm of the light transmission I o / I vs time, resulted in straight lines for all cases, establishing the first-
k = 1025.64 exp[.-- (79.5:l:2.5kcal/mole)/RT~ X cm3/mole" sec. This result is in rough agreement with data of other authors 2~ even though the evidence for the unimolecular character is much more cogent here, because of the wider range of conditions. The appearance of N H radicals during the decomposition of NHa at all temperatures above 2100~ suggests that these radicals play a major role in the mechanism.
362
REACTION
9.5
--
KINETICS
Reaction (1) as the primary decomposition mode of hydrazine and the first consecutive step (2) seem to be well established by these and other investigations 2.s:
*"
o o
,9.0
ee 9 .,r
NH2 + N2I-I4~ N~Ha + NHs.
(2)
.\
N~H~ -~ H + HNNH
\
~.5--
(1)
The chain mechanism, evidenced by the experimental results at lower temperatures, might include as the first propagating reactions:
~
tog k
N2H4 --~ 2 NH2
(3)
or
/ NH2 + N2H3 ~ NH3 + HNNH
NN
\ II
f
crn 3
7 9
7.5--d[Nt'r k fNH 7 [,4d I
0.35
I
(30
o~"\
NH~ + HNNH --. NHs + N2 + H.
,\kkk~ [
0.40 0.45 lIT.103(~ -'1
F m . 9. Arrhenius plot of second-order rate constants in NHs pyrolysis at temperature from 2100~
(5)
Combination of (3') and (5), excluding (3), would result in an ammonia production which is too high, even when further reasonable propagating or terminating reactions are taken into consideration, Spontaneous decomposition of N2H2 to N2 and H2 ['conceivable, if (3') forms the species
to 2900~ Discussion
In assigning a scheme of elementary reaction steps to the hydrazine decomposition under the described conditions, the following empirical facts have to be accounted for: (a) As one approaches higher temperatures and lower concentrations, the N2H4 decomposition becomes first-order in character. Elsewhere there seems to be an appreciable contribution of a chain mechanism, which gains in significance with increasing N2H4 content and gives rise to a pronounced induction period. (b) The ammonia produced at the lower temperatures corresponds closely to v = 1. With increasing temperature, v becomes smaller. At a reaction temperature of 2000~ no NHs could be detected with the described procedure and under the prevailing pressure conditions, v does not seem to depend very much upon the initial N2H4 concentration. A detailed discussion of possible elementary reactions is given elsewhere. 1~ Here, only the essential points are presented. The occurrence of
N N IH\ H
would not permit a reaction chain to develop. Hence (3) might be the dominant destruction reaction of N2H3, followed by H + N~H4 --* NHs + NH2
(4)
H + N2H4 ~ H2 -4- N2Ha
(4')
and the recombination reactions NH2 + NH~ -~ NH3 + NH
(6)
NH2 + NH2 --~ HNNH W H2
(6')
[M] + NH2 + H ~ NHs + [M].
(7)
If kt < k4f, practically no ammonia would be formed under conditions which favor chain decomposition (extrapolation of room temperature data 22,23to shock-wave conditions might not be quite iustified). By comparison with the H elimination from ethyl radicals, 24 the activation energy of (3) is estimated to be of the order of 42 keal/mole, v/z., (3) proceeds comparatively fast.
THERMAL DECOMPOSITION OF HYDRAZINE
Assuming that the measured k~' = kl and that chain terminating steps are not significant in the initial reaction phases, which are controlled only b y Reactions (1), (2), (3), and (4), an application of the steady-state approximations to the concentrations of N~H~ and H (by fitting the computed expression for the half-lives to experimental data) yields k~ ~ 1018'5exp [--- (17 kcal/mole)/RT~ X cm3/mole "sec. The theoretical curves derived with the above assumptions are shown in Fig. 2. An alternative interpretation on the basis of Reactions (1), (2), (3), and (4) might involve k2 ~ 10la cm3/mole-sec, ki' = 3kl and Reaction (3) being rate determining for the development of a chain mechanism. Numerical integration* of the four differential equations, however, has shown that this interpretation yields functions of half-lives, with respect to concentration and temperature, different from those observed experimentally. With kl' = kl and the above value for k2, scarcely any NH3 would be produced in chain propagating steps at temperatures above 1400~ In order to explain the experimental result v 0.5 at 1600~ with further decrease of v towards higher temperatures, the radical recombination (6'), followed by (5) is suggested. At temperatures above 2000~ Reaction (5) might be dwarfed by the unimolecular decomposition of H N N H HNNH --o NNH + H,
(8)
a reaction which m i g h t - - b y comparison with the NH3 decomposition--be expected to show strong dependence on total gas density, and which reduces the NH3 yields in the over-all decomposition. An appreciable contribution of Reaction (6) appears to be unlikely, because in that case the production of NH3 should be insensitive to variations in temperature. ACKNOWLEDGMENTS
We thank Professor W. Jost for his continuous interest and guidance in these investigations and the OAR for financial support through the European Office under Contract AF 61 (514)-1142. REFERENCES 1. AUDRIETH, L. F. AND OGG, B. A.: The Chemistry of ttydrazine: Wiley, 1951.
2. SZWARC, M.: Proc. Roy. Soc. (London) A198, 267 (1949). * The efficient cooperation of H. Eberius in solving this problem on an IBM 650 is gratefully acknowledged.
363
3. GRAY, P. AND LEE, I. C.: Seventh Symposium (International) on Combustion, p. 61, Butterworths, 1959. 4. HVSAIN, D. AND NORRISH, R. G. W.: Proc. Roy. Soc. (London) A273, 145 (1963). 5. GRAY, P., LEE, J. C., AND SPENCER, M.:
Combust. Flame 7, 320 (1963). 6. MOBERLu W. M.: J. Phys. Chem. 66, 366 (1962). 7. McHALE, E. T., KNOX, B. E., A~D PALMER, H. B.: Personal communication. 8. DIESEN, R. W.: J. Chem. Phys. 39, 2121 (1963). 9. MICHEL, K. W. AND WAGNER, I~. GG.: in Investigation of Gaseous Detonations and Shock-Wave Experiments with Hydrazine. Contract No. AF 61 (514)-1142, Technical Summary Report No. 2, 1960; MICHEL, K. W. AND WAGNER, H. GG. :
The Pyrolysis and Oxidation of Hydrazine in Shock Waves. Contract No. AF 61 (514)-1142, Technical Summary Report No. 3, 1962; MICHEL, K. W. AND WAGNER, H. GG.: in
Detonation and Shock-Tube Studies of Hydrazine and Nitrous Oxide. Contract No. AF 61 (514)1142, Annual Report No. 4, 1963. 10. MICHEL, K. W. AND WAGNER, H. GG.: Ber. der
Bunsen-Gesellschaft, in press. 11. MICHEL, K. W.: Ph.D. dissertation, G6ttingen, 1962. 12. RUDINGER, G.: Phys. Fluids $, 1463 (1961). 13. MICHEL, K. W. AND WAGNER, I-~. GG.: Z.
physik. Chemie NF 35, 392 (1962). 14. FONER, S. N. AND HUDSON, R. L.: J. Chem. Phys. 29, 442 (1958). 15. BURTLE, J. G.: Ind. Eng. Chem. ~4[, 1675 (1952). 16. MICHEL, K. W., OLSCHEWSKI, H. A., RICHTERINO, H., AND WAGNER, H. GG.: Z. physik. Chemie NF, in press. 17. Tables of Thermal Properties of Gases, National Bureau of Standards Circular 564, Washington, 1955. 18. LANDOLT--B6RNSTEIN: Kalorische ZustandsgrSssen, Vol. 2, Chap. 4, Springer, 1961. 19. BENNETT, m. G. AND DALBY, F. W.: J. Chem. Phys. 82, 1716 (1960). 20. MATHEWS, J. C., GIBBS, M. E., AND I~OLSEN,
J. N. : 139th National Meeting of the American Chemical Society, St. Louis, 1961. 21. JACOBS, T. A.: J. Phys. Chem. 67, 665 (1963). 22. BIRSE, E. A. B. AND MELVILLE, I-I. W.: Proc. Roy. Soc. (London) A175, 164 (1940). 23. SCHIAVELLO, M. AND VOLPI, G. G.: J. Chem. Phys. 37, 1510 (1962). 24. BrWATER, S. AND ROBERTS, R.: J. Chem. Phys. 19, 326 (1951). 25. DARWENT, D. DE B. AND ROBERTS, R.: Discussions Faraday Soc. 14, 55 (1953).
364
REACTION KINETICS
COMMENTS Prof. H. B. Palmer (Pennsylvania State University): Michel and I have had stimulating discussions about the compatibility of our studies, and it seems appropriate to mention what seem to me to be the critical points. The oscilloscope records of Michel and Wagner show induction periods t h a t are surely real. Their lengths correspond quite well to estimates of the times required to double the initial rate in the nonchain mechanism of our paper. I t is not obvious to me that, in their records, the rate increases by more than a factor of two during the induction period. If it really does, then our rejection of a chain must be wrong and we have to look for a source of error. The most obvious parameter to examine is the temperature in the reflected shock. There are two main questions a b o u t this:
temperatures above 1250~ Here, their experiments on the effect of the N2H4 concentration upon the half-life seem to show very beautifully the transition from a steady state for NH2 to a nonsteady state, i.e., from an over-all rate equal to twice the rate of the first step to a rate equal to one times it. A final point of interest on the mechanism is t h a t the nonchaln mechanism resembles t h a t for decomposition of hydrogen peroxide (Baldwin, 1%.,Jackson, D., Walker, R. W., and Webster, S. J. : this Symposinm p. 423). The parallels are: H202 --* 2 OH OH + H202 ~ H20 + HO2
N2H2 --* 2 NH2 NH~ + N2H~ --~ NH3 -t- N2H3
(a) How accurately known is the temperature of the initially-reflected shock? (b) How, if a t all, does the temperature in the reflected shock change with time?
The difference is:
In both shock-tube studies, ideal reflection is assumed. Although the shock tubes are somewhat different in geometry, errors from this assumption should be roughly the same in b o t h cases. Question (b) is, however, answered differently. Mainly by calculations, Michel concludes t h a t the temperature changes very little with time. From pressure records, we conclude t h a t it increases substantially during our dwell times. The pressure-increase phenomenon has been observed in at least four other laboratories and we are confident t h a t it is real. For example, it can be removed by modifying the flow in the shock tube [Tsang, W.: J. Chem. Phys. 40, 1498 (1964)]. When the pressure increase occurs, the gas must be heated in very nearly an isentropic manner. We do not find fault with Michel's calculations, b u t take rather the position t h a t it may not be correct to treat the reflected gas as a stagnant body adjacent to a cool wall (except at the endplate). We think t h a t aerodynamic effects are probably very significant in determining the temperature history. At short times, these temperature changes will be slight, which is why we have compared our data at low temperatures with Michel and Wagner's at
This makes us wonder whether due consideration has been given to the reaction OH q- HO2 -+ H~O + O2 as a possible third step in the decomposition of H202. I t gives the same kinetics as 2 HO2 ~ H202 + O.~, and should be very fast. Dr. K. W. Michel: I n Fig. 1, the rate of the decomposition of hydrazine increases b y more t h a n a factor of 2. I n general, absolute values of the initial reflected shock temperatures are not claimed to be more accurate, in this case, t h a n to about 30~ even though the consistency of relative temperatures which is material for the accuracy of the activation energy) is much better. This is also borne out by the excellent agreement of data obtained in different shock tubes of this investigation. The temperature increase of the reflected gas, due to adiabatic postcompression, gains in significance at reaction times in excess of 500 usec. Again, this is substantiated experimentally b y the observation of first-order rate laws over tong reaction periods in the decompositions of NH~, COs (Ref. 16) and N20 [JosT, W., MICHEL, K. W., TROE, J., AND WAGNER, H. Go.: Z. Naturforschung 19a, 59 (1964)] with the same equipment.
2 HO2 -o H202 + O~
N2H~ ~ NH~ --* NH~ -t- (NzH2) --+ N2 ~ H~.