1182
G . A. Takacs and G. P. Glass
Conclusions The results presented above show that most of the electrons produced by y-irradiation of liquid xenon a t 164°K behave as free electrons and are homogeneously distributed throughout the liquid which hllows the use of simple homogeneous kinetics in contrast to the complex kinetics used in liquid hydrocarbons.5 The yield of these free electrons, Gel(@-)= 4.3, is higher than that obtained for liquid argong where Gfi(e-) = 2.0. This observation with xenon provides further evidence for the increase of Gpi with the electron mobility in dielectric liquids and is in line with previous results, e.g., G f l = 0.13,3 we = 0.09 cm2 V - I sec-1;6 Gel = 0.332,3 we = 7 cm2 V-1 sec-1;6 Gf, = 0.857,3 we = 55 cm2 V - I sec-1;6 Gfl = 2.O,9 ye = 475 cm2 V-I sec-l;1° and Gel ='4.3, we = 2200 cm2 V-1 sec-l,lO for n-hexane, 2,2,4-trimethylpentane, neopentane, argon, and xenon, respectively. This yield of electrons of G(e) =
4.3 should be compared to that obtained in the gas phase G(e) = 4.621 and indicates, as in the case of hydrocarb o n ~ that , ~ the yields of electrons ( i e . , the W values) are similar in both phases. The product formation is understood in terms of the electron scavenging by N2O to give NzO- and the secondary reaction of the latter to give (NzO)O-. This ion recombines or yields one excess nitrogen by further reacting with N2O. The resulting anion ( ( N 2 0 ) 2 0 - ) reacts in competition with its recombination with two nitrous oxide molecules to give two excess nitrogen molecules. While the reaction of NzO- with nitrous oxide is thought to proceed with a diffusion-controlled rate, that of (N2O)O- is only 8.3 x lo4 1M-l sec-l. The three-body reaction involving (N20)zO- is estimated to be 4.7 x 103 M - 2 sec-1. Finally, the ratio of the rate constant for reaction of NzO- with CO2 to that for reaction with NzO is 2.5 indicating that the reaction with COz is also diffusion controlled.
Reaction of Atomic Oxygen with Hydrogen Bromide G . A. Takacs and G. P. Glass* Department of Chemistry, Rice University, Houston, Texas 77001 (Received January 8, 1973) Publication costs assisted by the Petroleum Research Fund
The reaction of atomic oxygen with HBr was studied in a fast discharge flow system by monitoring epr spectra of O(3P2), BrPPs/z), OH(2~3,2),and H(2&/2) a t various reaction times. The experimental results were found to be in accord with the mechanism, 0 HBr OH Br (l),OH HBr HzO + Br (2), OH + 0 H 0 2 (3), H HBr Hz Br (4), and k l was measured a t 298 K as 4.4 f 1.0 X 10-14 cm3 molecule-1 sec-1. The radical intermediate, OH(Za3/2), was detected in both the ground state, and in the first vibrationally excited state. Measurements on the relative population of these states allowed a lower limit of 0.3 to be placed on the ratio k(lbl/k(b), 0 HBr OH(v = 0) + Br(2P3/~) (la) and 0 HBr OH(u = 1) + Br(2P312)( l b ) . +
+
+
+
-
+
+
+
+
+
+
-
-
+
Introduction This paper reports part of a general study of reactions of hydrogen halides with simple free radicals ( e . g . , 0, H, and OH).1 Such reactions play an important role in the inhibition of flame and combustion processes.2,3 No previous direct experimental study has been made of the reaction of atomic oxygen with HBr. However, the reaction has been investigated theoretically by M a y a and Schieler4 using the BEBO method introduced by Johnston and Parr. They calculated the rate constant for reaction 1 as 3.15 X 10-12 ! P I 2 exp(-300/RT) cm3 molecule-I sec-1. Only one other estimate of this rate constant has appeared in the literature. In a study of the inhibition of the second explosion limit of the Hz-02 reaction by HBr, Clark, et al.,5 simulated the course of the reaction by numerical integration of the rate equations pertaining to a mechanism that included reaction 1. A reasonable fit to experiment a t 773 K was obtained when k l was set equal The Journei of Physical Chemistry, Vol. 77, No. 9 , 1973
to 4.15 X 10-11 cm3 molecule-1 sec-1. However, the fit was not sensitive to the actual value of k l if chosen within large limits.
0
+
HBr
-
OH
+
Br
(1)
In this study the reaction was followed directly in a fast flow system using epr detection. Spectra of O(3P2), OH(u = 01, OH(v = I),' H(2S1/z),. Br(2P3/~),and Br(2P1/2)7 (1) G. A. Takacs and G. P. Glass, J. Phys. Chem., accepted for publication. (2) R. N. Butlin and R . F. Simmons, Combust. Flame, 12, 447 (1968). (3) W. E. Wilson, Jr., J. T. O'Donovan, and R. M. Fristrom, Symp. Combust. 12th, 929 (1969). (4) S. W. Mayer and L. Schieler, J . Phys. Chem., 72, 236 (1968). (5) D. R. Clark, R. F. Simmons, and D. A. Smith, Trans. Faraday Soc., 66, 1423 (1970). (6) P. N. Clough, A. H. Curran, and B. A. Thrush, Chem. Phys. Lett., 7, 86 (1970). (7) P. B. Davies, B. A. Thrush, A. J. Stone, and F. D. Wayne, Chem. Phys. Lett., to be submitted for publication.
1183
Reaction of Atomic Oxygen with Hydrogen Bromide were recorded. An attempt was made to identify the primary reaction from the following possible choices
+
-
0 HBr OH(v = 0) -I- Br(2P,,,) AH = -14.8 kcal/mol (la) 0 HBr OH(v = 1) + Br(2P3,z) A H = -4.6 kcal/mol (lb) 0 HBr -,OH(v = 0) Br('Pljz) AH = -4.3 kcal/mol (IC)
+
+
+
+
Experimental Section
The construction and operation of the discharge flow system and epr spectrometer have been described previously.l,S Atomic oxygen was produced by either microwave discharge of 02-Ar mixtures, or, in the absence of 0 2 , by NO titration of atomic nitrogen9 N + N O -
N2+O
HBr was added to the 20-mm i.d. quartz flow tube through a small movable inlet probe. Concentrations of 0, Br, H, and OH were measured absolutely by epr using the 0 2 reference method. The initial concentration of HBr was estimated from its measured flow rate. Prior to each set of experiments, the flow tube was treated with a fluorinated halocarbon coatinglo that has previously been shown to inhibit wall recombination of atomic bromine and many other free-radicals species.l The corrosive reaction products were trapped a t 77 K in a U-tube positioned immediately upstream of the pump. A large brass dump tank placed between the flow tube and the liquid nitrogen trap effectively removed all trace of ozone. In the absence of this tank, a highly explosive mixture of ozone and reaction products was condensed. For this study the epr system was slightly modified to improve its sensitivity. A new cylindrical cavity, operating in the TEolz mode, and having end plates that can be screwed in or out to change its length, was constructed. Improved sensitivity was obtained by adjusting the end plates until the position of maximum microwave field was superimposed on that of maximum modulating magnetic field. With this cavity OH concentrations of lof1 molec u l e / ~ ~ $could be detected, although accurate quantitative measurements were not possible at this concentration. Results
In this study the growth of Br(2P312) and the decay of atomic oxygen was followed as a function of reaction time in 12 different mixtures for which the ratio of the initial concentration of HBr to that of atomic oxygen varied from 2.02 to 34.6. The concentration of OH and of atomic hydrogen was also monitored in several of the more stoichiometric mixtures. The results of these measurements were classified and tabulated using the following reaction mechanism
-
0 + HBr OH + Br (1) OH + HBr -r HzO + Br (2) OH+O O,+H (3) H + HBr H, + Br (4) Reactions 2, 3, and 4 are well known, and their rate constants have been measured as 5.1 X 10-12, 4.3 X 10-11, and 3.4 x cm3 molecule-1 sec-1, respectively.l.8
-L
The alternative reactions of atomic oxygen with HBr, namely
+
-
HBr H + BrO (5) 0 + HBr HOBr (6) were dismissed since reaction 5 is 31 kcal/mol endothermic, and reaction 6 produces a 3 atom complex whose lifetime is too short for it to be collisionally deactivated. Reactions 7 and 8 were neglected because they are too slow to be of significance in the presence of atomic OXYgen. Under the present conditions, atomic bromine should be unreactive since wall recombination is effectively inhibited by the fluorinated halocarbon wall coating, thirdorder recombination is slow at pressures below 1 Torr, and the abstraction reactions with HBr, 02, and HZ are ver:y endothermic. OH + OH HzO + 0 (7)
0
. )
OH
--
wall
(8)
The correctness of the mechanism was tested by ( i ) comparing the kinetic data on oxygen atom decay and Br(2P3/2) growth with that predicted by the mechanism, (ii) determining the concentration of the predicted intermediates (H and OH), and (iii) measuring the reaction stoichiometry. An analysis of the kinetic behavior predicted by the mechanism is straight forward only when HBr is present in large excess. In this limiting case, all OH reacts via reaction 2, and reactions 3 and 4 can be neglected. Since the concentration of excess HBr remains effectively conistant throughout the reaction, atomic oxygen undergoes pseudo-first-order decay, and its concentration at any time t is given by the equation In (O),/(O) = kl(HBr)t
(9)
where ( 0 ) o represents the initial oxygen atom concentration. If 122 >> kl, reaction 1 will be rapidly followed biy reaction 2, and two bromine atoms will be produced for each oxygen atom reacted. Thus (Br) = 2[(0)0 - (O)], (Br)m= 2(0)0, and In (Br)-/[(Br)-
-
(Br)] = k,(HBr)t
(10)
Equations 9 and 10 predict that plots of In (0) or In [(Br)m - (Br)] against reaction time should be linear, and have identical slopes of magnitude kl(HBr). Figure 1 illustrates such plots for two reaction mixtures: one with (HBr)o/(O)o = 34.6, and the other with (HBr)o/(O)o = 15.5. For these mixtures (Br)- was estimated as the bromine atom concentration determined at the longest measured reaction time plus twice the atomic oxygen conceritration remaining at that time. Values of k l estimated from Figure 1 and other similar plots are tabulated in Table I. Also included in this table are values of k l computed from experimental data, assuming the reaction mechanism to be composed of reactions 1-4. The computed values are chosen to minimiz,e the difference between the experimentally measured 0x37gen atom concentrations and those calculated by integrating the rate equations pertaining to reactions 1-4. A sirnple Newtonian integration method was used with a step size of sec. In order to avoid the mixing zone, i n k (8) J. E. Ereen and G. P.Glass, J . Chem. Phys., 52, 1082 (1970). (9) F. Kaufman, J. Chem. Phys., 28,992 (1958). (10) Coating supplied by Marchem Inc., P. 0.Box 6914, Houston, Tex. 77005. The properties of this coating are fully described in ref 1. The Journal of Physical Chemistry, Voi. 77, No. 9, 1973
1184
G . A. Takacs and G. P. Glass
TABLE I: Kinetics of the Reaction of Atomic Oxygen with HBr
(O)O>
(Ar), 10’6 molecule cc-1
(HBr)o, 10’4 molecule cc-1
1014 molecule cc-1
0 decay
1.24 1.95 2.33 2.36 2.17 2.32 2.04 2.02 2.21 3.20 2.64 1.a3
3.84 11.3 22.7 21.3 18.5 17.1 13.4 26.7 19.6 29.7 25.4 35.4
1.71 2.10 3.84 2.40 1.75 1.36 0.9Ob 1.72 0.97b 1.256 0.986 1.026
5.9 3.6 3.7 5.1 3.7 6.6 5.7 5.7 6.4 7.1 5.3
a k(calcd)
= 4.4 f 1.O X 10-
l4cm3 molecule-
sec-
k,
cc molecule-’ sec-’ Br growth
Calcda
5.7 6.3 7.5 5.4
4.9 2.6 2.8 3.8 3.6 5.5 4.6 4.8 5.7 5.6 4.4
(Br formed/ 0 reacted)
(Br formed/ 0 reacted),,i,d
1.15 1.61 1.55 1.62 1.91 1.52
1.24 1.35 1.45 1.52 1.60 1.61
1.71 1.76 1.65 2.00 1.98
1.73 1.74 1.76 1.a2 1.aa
* For these mixtures 0 was produced in the absence of 02 by NO titration. reaction 1. When this occurs the rate constant obtained by pseudo-first-order oxygen decay (graphically) is greater than kl. Since k3/k2 = 8.4, even a mixture with an 8.4fold excess of HBr reacts to form (initially) equal amounts of H2O and 0 2 , and with a pseudo-first-order rate constant (measured using the early points) 50% greater than ki. A successful search was made for the two intermediates, OH and H, predicted by the mechanism. Unfortunately, due to their low concentration they could be detected only a t short reaction times and only using stoichiometric mixtures. Under these conditions their spectra were observed with SIN ratios of between 2 and 5 , and only rough estimates could be made of their absolute concentrations. These estimates are recorded in Table I1 together with values of (OH) and (H) that were calculated from the mechanism using the value of kl measured in this study, namely, 4.4 X lO-I4 cm3 molecule-1 sec-I, and the,relationships (OH) = h,(O)(HBr)/k,(HBr)
‘2
4
6 8 IO TIME ( m s e c )
12
14
16
Plot of log (0) [open circles and squares] and log ((Br)” - (Br)) [closed circles and squares] vs. reaction time: (a) (HBr)o = 3.54 X 10l5, (O)a = 1.02 X l o q 4molecule/cc; (b) ( H B r ) o = 2.67 X ( 0 ) o = 1.72 X l o i 4 molecule/cc.
Figure 1.
gration was begun at the first measured point. The rate constants 122, k3, and k4 were not varied during computation, and were chosen as 4.3 X 10-11, 5.1 X 10-12, and 3.4 X 10-12 cm3 molecule-1 sec-1, respectively.1J k1 was varied in increments of 10-15 cm3 molecule-1 sec-1 until the best fit between experiment and calculation was obtained. In order to reduce computation time the initial trial value of k1 was set equal to the value estimated from Figure 1. It is clear from Table I that, in general, the computed value of kl is less than the value obtained by pseudo-firstorder oxygen atom decay. This is to be expected since the overall reaction stoichiometry changes with the (HBr)/(O) ratio, being governed by the relative rates of reactions 2 and 3. When HBr is in great excess, all OH is consumed by reaction 2, and thus the rate of removal of 0 is equal to the rate of reaction 1. If more atomic oxygen is present, some OH is consumed in reaction 3, and the rate of removal of atomic oxygen is then greater than the rate of The Journal of Physical Chemistry, Vol. 77, No. 9, 1973
+ h3(0)
(11)
and which were derived using the steady-state approximation. This approximation can be applied to OH since kl
