- Email: [email protected]
Laser photolysis of ozone in the presence of ammonia. Formation and decay of vibrationally excited NH2 radicals
Valume 120, number 1
CHEtilCAL
27 September
PHYSICS LETTERS
i985,
LASER PHOTOLYSIS OF OZONE IN THE PRESENCE OF AMMONIA. FORMATION AND DECAY OF VIBRATIONALLY EXCITED NH, RADICALS S-G. CHESKIS,
A.A.
IOGANSEN,
0-M.
SARKISOV
fmrirure of Chemical Physics. USSR Academy of Sciences, I I7334
and A.A. Moxcon:
TITOV
USSR
Received 24 June 1985
NH, radicals were generaled by laser pulsed phoIolysis or O3 in the presence or NH, and their subsqurnr kinetics were monitored by LIF. The room-temperalure rate consLant for O(‘D)+NH, WIS measured IO be (3.3_~0.1)~10~‘” cmm3 molecule-’ 5-l by sludying NH2 accumulation profiles. The nascent energy dislribulion over NH, vibrational levels was studied and the fraction or NH, radicals generated in the ground vibralional srare was round 10 be 0.42~0.08. Acceleratior. oi the reaclion NH2 +O, caused by vibrational exciration or NHI radicals was observed - the rare consmnfi were (lSf0.4)~ IO-” and (1.5~0.3)~10-‘~ cm-’ molecule-’ 5-l. [or vibrationally exciled and non-exciled NH2. rcspeclively. The reactive channel lor removal of vibralionnlly excited NH, by ozone appears LObe dominant.
1. Introduction The chemical reactions of excited species are of current interest for a number of reasons. Thus, the elucidation of chemical mechanisms of various nonequilibrium processes (chemical lasers, plasmas, etc.) requires precise quantitative information about stateto-state reaction rate constants. On the other hand, these constants and their dependence on the form of available energy excess may help to reveal the detailed dynamics of an elementary chemical collision. The pulsed laser ozone photolysis in the Hartley band is a convenient laboratory system for such investigations_ The metastable oxygen atoms O(lD), produced in a short pulse via O3 + hv + O(‘D) + 02( ‘A)
(1)
usually react rapidly with an added gas (hydro-, halocarbons, hydrides, etc.). Due to their high exoergicity, many O(lD) reactions are intense sources of excited particles with a definite energy distribution Here we report results of the LIF monitoring of ground-state NH; kinetics after pulsed laser ozone photolysis in the presence of N&,_ The reaction 0(‘D)+NH3’+N&+0,H 0 009-2614/85/S (North-Holland
(2)
03.30 Q l$evier Science Publishers B-V. physics Publishing Division)
and subsequent reactions of the vibrationally excited and ground-state NH, radicals have been investigated. The OH(u) temporal behavior in this system will be the subject of a later publication.
2.Experimental The apparatus and procedures were described in ref. [I]. The ozone was photolysed by 10 ns pulses of a FQ Nd3+ r YAG laser at 266 nm (l-2 mJ/pulse). The laser beam was unfocused to reduce the diffusion decay of radicals. The LIF technique was used to monitor the concentration of NH2 in the ground state. A prism-narrowed (AX = 1 A) R6G dye laser, pumped by a FD Nd* :YAG laser, was tuned to the strong absorption in the NH, spectrum corresponding to the (0,9,0) vibronic band at 5979 _&.The fluorescence was collected by af/l.S lens system and. after trans mission through the absorption cut-off falters (KS1 I), was detected by a F. 136 photomultiplier with a 1 cm diametz S20 cathode at room temperature. The electric signal was strobed, amplified (X 100) and stored in digital form. An on-line microcomputer was used for data storage and processing The time delay between photolysis and probe pulses was also 45
Volume 120. number 1
microcomputer controlled. Usually 20-40 pulses were averaged for each time delay. Ammonia and ozone (diluted in Ar) were mixed in a slow flow before entering the fluorescence cell (black anodized Al waUs). Ozone was prepared in the low-frequency electric discharge from 02. All the gases were purified by successive freeze--thaw cycles under vacuum. The total gas pressure was monitored by a diaphragm manometer with an accuracy of 0.05 Torr. The inital O(lD) concentration was estimated to be lower than 1013 cm-3 from the photolysis beam fluence (1014 photons cm-*) and ozone cross section.
3. Results
and discussion
2 1. Rate comtwlt of the reaction O[lD) + NH3 So far the rate constant of reaction (2) has been measured only by studying d(lD) decay kinetics [2,3]. Fletcher and Husain [2] monitored O(lD) concentration using VUV absorption in the 115 nm region. Davidson et al. [3] observed a weak emission corresponding to the forbidden transition 0(2 JD) + 0(2 3P2). Reported k, values, 6.7 X lo-lo [2] and 2.5 X lo-Jo cm3 s-l [3], are in poor agreement; moreover, the discrepancy has a common character for a variety of o(JD) reactions [3] _ In this work a different method was used for k2 determination - the monitoring of the NH2 accumulation profile. The kinetics of reaction (2) &as investigated under the following conditions: [03] = 1 X 1014, [Ar] = 3 X 1016, and [NH31 varied in the range 8 X 10131.3 X 1015 molecule cm-g_ Under these conditions, the time constant of NH2 accumulation, and that of O(lD) decay, ranged from 2 to 20 ps Only three reactions may appreciably affect the NH2 concentration profile on this timescale, the reaction (2) and O(lD) quenching ractions: O(lD)
+ 03 + products,
0(1D)+Ar+O(3P)+Ar.
If only centration
NW,
reactions (2)-(4) take place, at a time t is given by
= W-&P
- exp(--k,dl
E
1985 [
the NH2 con-
where ken = k2 [NH31 + k3 [03] + k4 [AI]. is the overall pseudo-first-order rate constant for O(lD) decay and [NH21 o is the maximum concentration of NH2 formed in reaction (2) in the ground state. From the experimental data, (I) gives the value of lierr as the slope of a linear plot of ln(l - [NH2],/ [NH*] o) versus t. The quantity k,~ was found to be linearly dependent on [NH31 as shown in fig l_ The bimolecular rate constant of reaction (2) was obtained from an unweighted linear least-squares slope of the plot in fig. 1 as k2 = (3.3 rf: 0.1) X IO-JO cm3 s-l_ The value of k, agrees within the error limits with that reported by Davidson et al [3], but is lower by a factor of 2 than the result obtained by Fletcher and Husain 123. Such a discrepancy may be ascribed [3] to the uncertaintly in the factor 7 used in [2] in a modified Beer’s law: 1= I, exp[-e(cZ)7] _ 3.2. The effect of uibrationalexcitation of NH, on NH2 + 0, reaction rate The reaction NH2 -to3 AH$&
-NH20+02, = -35
kcal/mole ,
(5)
(4) 0
5
10
f
I (1) i i
,
(3)
All other processes appear to be unimportant. Thus, the fastest competing process, the NH2 vibrational relaxation, proceeds about 30 times slower than these reactions. 46
27 September
CHEMICAL PHYSICS LElTERS
is
Fig. L Pseudo-first-order rate constant for O(‘D) decay ke versus ammonia concentration Pmt = 6 Torr.
I 1 1 ]
I
Volume
120,
number 1
CHEMICAL
PHYSICS
has been investigated directly only in the last five years [1,4-91. The rate constants at room temperature reported in refs. [ 1,5-71 agree fairly well, and lie in the range (1.3-1.9) X lo-l3 cm3 s--r_ Kurasawa and Lesclaux [4] obtained about half thisvalue wit&k, ~(0.63 f 0.1) X lO-l3, while.in the recent work by Patrick and Golden [S] kg was found to be 3.3 X lo-13 cm3 s-l_ On the basis of kg temperature dependence data [4,6-S], the value of 2 to 3 kcal/mole may be recommended now for the activation energy of reaction (5). However, what particular form of excess energy is more effective in overcoming the activation energy barrier is a question that is to be investigated_ There are few reliable data concerning the effect of the vibrational excitation of radicals on their reaction rates in the literature_ In this work we have measured an acceleration of the reaction (5) caused by vibrational excitation of NH2 radicals. The experiments were carried out under the following conditions: [Ar] = 1.8 X 1017; [NH31 = 5 X 1Ol4 and 1 X 1015; [03] was varied from 9 X 1014 to 1.1 X 1 016 molecule cmm3_ Under these condition% the O(lD) atom lifetime did not exceed 2 F. We monitored the NH* kinetics in the time interval 6~s to 2 ms after the photolysis pulse. On this timescale the major processes which contribute to the NH, temporal profde, apart from reaction (5) are the following *: NH;+Ar+NH*+Ar,
(6)
NH;+NH3+NH,+NH3,
(7)
NH’2+03-+NH,+03,
(8)
NH; •i-O3 + products _
(9
LETTERS
27 September
it proceeds mainly via physical quenching [lo]. Moreover, we carried out control runs varying the photolysis beam fluence and hence changing initial concentrations of active species. No noticeable changes were observed in the NH2 kinetics, providing additional evidence for the postulated mechanism (5)-(9)The solution of the differential equations corresponding to the scheme (S)-(9) gives for [NH,] in the ground state:
[NH21, = EN&l o exd--k#) + [NH;] 0 [kG/(kG +kc -krIl
where k, = k, [03], X-z = kg [03], k: = kg [Ar] + k7 [NH31 + kg 1031, and INHzl,-,and PHIIo are the initial concentrations of the ground-state and vibrationally excited NH, formed in reaction (2) In accordance with (II) the observed NH, kinetics had a “rise-and-fall” form as shown in fig 2. In the fall region (t > 300 us in our conditions) the observed curves are accurately exponential, giving the reaction (5) pseudo-first-order rate constant k,_ By varying the ozone concentration we have obtained the bimolecular rate constant of the ground state NH2 reaction z;h ozone: kg = (1.5 f 0.3) X lo-l3 cm3 molecule-l The much faster increasing portion of the curve corresponds to the NW2 decay processes (6)-(9). Using the known k, value, we may multiply the eq. (II) by exp(k,t), and so reduce it to a simple saturation curve:
0
0
I 0
8
*
1
1
400
Fig. 2 Ground-state NH2 kinetics Hereafter NH:
denotes viirationally excited
NH2
radi&
Icxp(--krt)
- exp [-(k: + k:) tl1,
Other active species, such as OH(u= 0, 1, 2, 3), 0(3P), O,(lA), were formed in our system along with NH; and NH,. However, it may be shown that they had negligible effect on the NH, kinetics The typical interaction time constants of 0(3P), OH and O&A) with NH2 and NH3, caculated from the reported rate constants and estimated initial concentrations, substantially exceed the measurement timescale. The interaction of vibrationally excited hydroxyl radicals with ammonia also may be neglected because *
1985
[O,]
TIME/&s [NH31
ml
= 9 X 1014 cmm3,
= 1 x 1016 cm-3,Pt&=6Torr.
47
CHEMICAL PHYSICS LEl-l-ERS
Volume 120, number 1
C(r)=C,-,+aCo[ l-
exp(-k&t)]
where C(f) - [NH21t exp(k,t) mental data, c= [I-3210>
co’= mH;lo,
,
(III)
is corrected
experi-
kEm=k;+k:-kFc,,
a=.k*g/k;KEq_ (III) may easily be transformed using semilog coordinates: In [ I - C(r)/(C,
= I,,
[ac; /(co
to a linear form
+ oCg*)] + c&,*)1 -
x-$-rr-
(W
It is to be noted that Co + crc*u represents a maximum level of the ground-state NH, concentration after the relaxation is over, and it may be determined from the corrected experimental curves (III) as the saturation level. The experimental data were plotted using semilogarithmic coordinates [In [I - C(t)/(Co + aCG)] , t] and the relaxation pseudo-first-order rate constant k& was obtained as a slope for each profile. As shown in fig. 3, k&was found to be linearly dependent on [03] through the whole probed [Ox] range, the upper and lower lines corresponding to 1 X 1015 and 5 X 1Or4 crne3 in NH, concentrationS respectively. Thus, the slopes of the lines in fig. 3 give, after correction fork,, the total second-order rate constant of the vibration&y excited NH, radicals removal by
ozone: kg + k8 = (1.5 f 0.4) X lo-r2 cm3 molecule-t s-1. From the difference between the intercepts in fig. 3, we estimate the rate constant of NH; quenching byNH3: k,=9X lo-l2 cm3 molecule-l s-r_ The good agreement of this k, value with Sarkisov et al. [ 1 l] probably-indicates that in our system, as well as in the ammonia photolysis system, the reaction (7) rate is limited by quenching of the low-lying bending mode levels of NH, radicals The rate constant of NH; removal by ozone is too high to be explained simply in terms of V-V transfer because the energy difference between vibrational levels of NH2 and O3 is large - 400 cm-t. So it seems to be of considerable interest to determine the relative importance of reactive and quenching channels. Therefore we have made an attempt to measure the k8/kg ratio. This ratio can be determined here because the relative increase in ground-state NH2 concentration due to the physical quenching depends on the percentage of NH; removed by reaction (9). If, for instance, all the NH; radicals took part in reaction (9), then there would be no increase in [NH2]. We can consider this effect quantitatively as follows: According to (IV), transforming the experimental data to linear form gives In [crC6/(Co + crCi)] as the intercept. The relative portion of NH2 concentration caused by NH; decay fl= oC~/(Cn + oG> may be written as l/P = 1 + Co/C;
= 1 +c,/c,
o-i +-2
0
,
,
4
8
[oJ/iO"cmm'~
~
I2
Fig. 3. Pseudo-&st+rder rate m&ant for vibrationally cxcited NH2 decay x-k versus ozone concentration Ptot = 6 Torr.1--[NHs]=l~1015cm-3,2-[NH~]=5X1014 cmm3_
48
+ (C,-,/C;)(k:
- k,)/k;
+(C&,)
~W~-k~)I0~lI(k~LW
+
’[
27 September 1985
+bINH31 +kg &I)(v)
If there is no vibrational acceleration of reaction (5) (i.e. if kg = k,),the value of P_l does not depend on the ozone concentration But if physical quenching is negligible, B_l grows linearly with [O,]. Fig. 4 shows the dependence of p1 on [03] obtained from our experiments. In spite of scatter in the data the increase of P_l with [03] is clearly visible. The slope of a least-squares line in fig. 4 gives the value of kg - kg, which leads to kg = 1.2 X IO-l2 cm3 s-l_ Thus, the reactive channel (9) is the dominant one and contributes about 80% to the NH; removal by ozone. The comparison of this kg value with the Arrhenius
I
CHEMICAL PHYSICS LETTERS
Volume 120; number 1
‘cl Fig 4. Dependence
I 5
fO-‘5-[0,, cm-’ 10
of p-r on ozone concentration
27 September
1985
conclusion that (35 + 15)s of the exothermicity of reaction (2) (that is 15 i 6 kcal/mole) is distributed among the NH, vibrational modes. This value is in disagreement with the present work. In facf assuming linearity of the suprisal plot (as ’ was done by Kinsey et al.), we estimate the mean vibrational energy of NHi to be 6 kcal/mole using the measured value of the ground-state NH, fraction, 42%. In our opinion, this discrepancy may be accounted for by errors in the energy balance method. Another explanation would be a strong bimodality in the NH, vibrational energy disposal_
(see text). References
expression for the rate constant of reaction (5), k5 = 4 X lo-l2 exp(-llOO/~ cm3 s-1 [7], shows that vibrational energy of NH, may be characterized by a 30% efficiency in overcoming the activation barrier. As yet we do not know which particular viirational mode and level correspond to the acceleration of reaction (5). 3. I lViYz nascent vibrational energy distn_butioIl Extrapolation to zero ozone concentration in fig 4 gives the value of 1 + Co/C,* = 1.7 f 0.2 (see expression (V)). Thus only (442 f 8)% of NH, formed in reaction (2) is in the ground vibrational state. This result may be compared with the data obtained for reaction (2) by Kinsey et al. [ 12]_ They determined the relative populations of u = 0 and u = 1 OH vibrational levels from the measured OH(O) and OH(u= 1) rotational energy disposal. In addition to this, they made approximate measurements of the mean translational energy of the products, and the mean rotational energy of NH,. Energy balance Ied the authors to the
[l]
A.A. logansen, 0-M. Sarkisuv and S.G. Chesk&, Khirn. Phyr (1985) 3-372 [in Russian]. I21 I..% Fletcher and D. Hunti. Can J_ Chem. 54 (1976) 1765. [31 J-A. Davidson, H-1. Schiff, G.E. Streit. J.R. hlcAfee, A.L. Schmeltekopf and CJ. Howard, J. Chem. Phyr 67 (1977) 5021. 141 H. Kurasawa and R. Lcsclaux. Chem Phyr Letters 72 (1980) 437. 151 V.P. Bulatov. A.A. Buloyan, S-G. Cheskis M-2. KozIincr, O.M. Sarkisov and A-1. Trostin, Chem. Phys Letters 74 (1980) 288. 161 W. Hack, 0. Horie and H.Gg. Wagner, Bcr. Dunserges Physik_ Chem 85 (1981) 72. 171 V.P. Buktov. AA. Buloyan, S.G. Cheskis, A.A. Iogansen and O.hl. Sarkosov, Khim Phys (1982) 2-513 [in Russian]. [8] R. Patrick and D-M. Goldeq J. Phyr Chem. 88 (1984) 491. [9] R. Lesclaw Rev. Chem Intermediates 5 (1984) 347. [IO] G.P. Glass H. Endo and B.K. Chaturvcdi J- Chem. Phyr 77 (1982) 5450. [ 1 l] S.G. Cheskh O.hl. Sarkosov and S.Ya Umzmsky, Dokl Akad Nauk SSSR 246 (1979) 66 [inRursian]. 1121 J.F. Cordova, CT. Rettner and J-L. Kinscy, J. Chem Phyr 75 (1981) 2742.
49
CHEtilCAL
27 September
PHYSICS LETTERS
i985,
LASER PHOTOLYSIS OF OZONE IN THE PRESENCE OF AMMONIA. FORMATION AND DECAY OF VIBRATIONALLY EXCITED NH, RADICALS S-G. CHESKIS,
A.A.
IOGANSEN,
0-M.
SARKISOV
fmrirure of Chemical Physics. USSR Academy of Sciences, I I7334
and A.A. Moxcon:
TITOV
USSR
Received 24 June 1985
NH, radicals were generaled by laser pulsed phoIolysis or O3 in the presence or NH, and their subsqurnr kinetics were monitored by LIF. The room-temperalure rate consLant for O(‘D)+NH, WIS measured IO be (3.3_~0.1)~10~‘” cmm3 molecule-’ 5-l by sludying NH2 accumulation profiles. The nascent energy dislribulion over NH, vibrational levels was studied and the fraction or NH, radicals generated in the ground vibralional srare was round 10 be 0.42~0.08. Acceleratior. oi the reaclion NH2 +O, caused by vibrational exciration or NHI radicals was observed - the rare consmnfi were (lSf0.4)~ IO-” and (1.5~0.3)~10-‘~ cm-’ molecule-’ 5-l. [or vibrationally exciled and non-exciled NH2. rcspeclively. The reactive channel lor removal of vibralionnlly excited NH, by ozone appears LObe dominant.
1. Introduction The chemical reactions of excited species are of current interest for a number of reasons. Thus, the elucidation of chemical mechanisms of various nonequilibrium processes (chemical lasers, plasmas, etc.) requires precise quantitative information about stateto-state reaction rate constants. On the other hand, these constants and their dependence on the form of available energy excess may help to reveal the detailed dynamics of an elementary chemical collision. The pulsed laser ozone photolysis in the Hartley band is a convenient laboratory system for such investigations_ The metastable oxygen atoms O(lD), produced in a short pulse via O3 + hv + O(‘D) + 02( ‘A)
(1)
usually react rapidly with an added gas (hydro-, halocarbons, hydrides, etc.). Due to their high exoergicity, many O(lD) reactions are intense sources of excited particles with a definite energy distribution Here we report results of the LIF monitoring of ground-state NH; kinetics after pulsed laser ozone photolysis in the presence of N&,_ The reaction 0(‘D)+NH3’+N&+0,H 0 009-2614/85/S (North-Holland
(2)
03.30 Q l$evier Science Publishers B-V. physics Publishing Division)
and subsequent reactions of the vibrationally excited and ground-state NH, radicals have been investigated. The OH(u) temporal behavior in this system will be the subject of a later publication.
2.Experimental The apparatus and procedures were described in ref. [I]. The ozone was photolysed by 10 ns pulses of a FQ Nd3+ r YAG laser at 266 nm (l-2 mJ/pulse). The laser beam was unfocused to reduce the diffusion decay of radicals. The LIF technique was used to monitor the concentration of NH2 in the ground state. A prism-narrowed (AX = 1 A) R6G dye laser, pumped by a FD Nd* :YAG laser, was tuned to the strong absorption in the NH, spectrum corresponding to the (0,9,0) vibronic band at 5979 _&.The fluorescence was collected by af/l.S lens system and. after trans mission through the absorption cut-off falters (KS1 I), was detected by a F. 136 photomultiplier with a 1 cm diametz S20 cathode at room temperature. The electric signal was strobed, amplified (X 100) and stored in digital form. An on-line microcomputer was used for data storage and processing The time delay between photolysis and probe pulses was also 45
Volume 120. number 1
microcomputer controlled. Usually 20-40 pulses were averaged for each time delay. Ammonia and ozone (diluted in Ar) were mixed in a slow flow before entering the fluorescence cell (black anodized Al waUs). Ozone was prepared in the low-frequency electric discharge from 02. All the gases were purified by successive freeze--thaw cycles under vacuum. The total gas pressure was monitored by a diaphragm manometer with an accuracy of 0.05 Torr. The inital O(lD) concentration was estimated to be lower than 1013 cm-3 from the photolysis beam fluence (1014 photons cm-*) and ozone cross section.
3. Results
and discussion
2 1. Rate comtwlt of the reaction O[lD) + NH3 So far the rate constant of reaction (2) has been measured only by studying d(lD) decay kinetics [2,3]. Fletcher and Husain [2] monitored O(lD) concentration using VUV absorption in the 115 nm region. Davidson et al. [3] observed a weak emission corresponding to the forbidden transition 0(2 JD) + 0(2 3P2). Reported k, values, 6.7 X lo-lo [2] and 2.5 X lo-Jo cm3 s-l [3], are in poor agreement; moreover, the discrepancy has a common character for a variety of o(JD) reactions [3] _ In this work a different method was used for k2 determination - the monitoring of the NH2 accumulation profile. The kinetics of reaction (2) &as investigated under the following conditions: [03] = 1 X 1014, [Ar] = 3 X 1016, and [NH31 varied in the range 8 X 10131.3 X 1015 molecule cm-g_ Under these conditions, the time constant of NH2 accumulation, and that of O(lD) decay, ranged from 2 to 20 ps Only three reactions may appreciably affect the NH2 concentration profile on this timescale, the reaction (2) and O(lD) quenching ractions: O(lD)
+ 03 + products,
0(1D)+Ar+O(3P)+Ar.
If only centration
NW,
reactions (2)-(4) take place, at a time t is given by
= W-&P
- exp(--k,dl
E
1985 [
the NH2 con-
where ken = k2 [NH31 + k3 [03] + k4 [AI]. is the overall pseudo-first-order rate constant for O(lD) decay and [NH21 o is the maximum concentration of NH2 formed in reaction (2) in the ground state. From the experimental data, (I) gives the value of lierr as the slope of a linear plot of ln(l - [NH2],/ [NH*] o) versus t. The quantity k,~ was found to be linearly dependent on [NH31 as shown in fig l_ The bimolecular rate constant of reaction (2) was obtained from an unweighted linear least-squares slope of the plot in fig. 1 as k2 = (3.3 rf: 0.1) X IO-JO cm3 s-l_ The value of k, agrees within the error limits with that reported by Davidson et al [3], but is lower by a factor of 2 than the result obtained by Fletcher and Husain 123. Such a discrepancy may be ascribed [3] to the uncertaintly in the factor 7 used in [2] in a modified Beer’s law: 1= I, exp[-e(cZ)7] _ 3.2. The effect of uibrationalexcitation of NH, on NH2 + 0, reaction rate The reaction NH2 -to3 AH$&
-NH20+02, = -35
kcal/mole ,
(5)
(4) 0
5
10
f
I (1) i i
,
(3)
All other processes appear to be unimportant. Thus, the fastest competing process, the NH2 vibrational relaxation, proceeds about 30 times slower than these reactions. 46
27 September
CHEMICAL PHYSICS LElTERS
is
Fig. L Pseudo-first-order rate constant for O(‘D) decay ke versus ammonia concentration Pmt = 6 Torr.
I 1 1 ]
I
Volume
120,
number 1
CHEMICAL
PHYSICS
has been investigated directly only in the last five years [1,4-91. The rate constants at room temperature reported in refs. [ 1,5-71 agree fairly well, and lie in the range (1.3-1.9) X lo-l3 cm3 s--r_ Kurasawa and Lesclaux [4] obtained about half thisvalue wit&k, ~(0.63 f 0.1) X lO-l3, while.in the recent work by Patrick and Golden [S] kg was found to be 3.3 X lo-13 cm3 s-l_ On the basis of kg temperature dependence data [4,6-S], the value of 2 to 3 kcal/mole may be recommended now for the activation energy of reaction (5). However, what particular form of excess energy is more effective in overcoming the activation energy barrier is a question that is to be investigated_ There are few reliable data concerning the effect of the vibrational excitation of radicals on their reaction rates in the literature_ In this work we have measured an acceleration of the reaction (5) caused by vibrational excitation of NH2 radicals. The experiments were carried out under the following conditions: [Ar] = 1.8 X 1017; [NH31 = 5 X 1Ol4 and 1 X 1015; [03] was varied from 9 X 1014 to 1.1 X 1 016 molecule cmm3_ Under these condition% the O(lD) atom lifetime did not exceed 2 F. We monitored the NH* kinetics in the time interval 6~s to 2 ms after the photolysis pulse. On this timescale the major processes which contribute to the NH, temporal profde, apart from reaction (5) are the following *: NH;+Ar+NH*+Ar,
(6)
NH;+NH3+NH,+NH3,
(7)
NH’2+03-+NH,+03,
(8)
NH; •i-O3 + products _
(9
LETTERS
27 September
it proceeds mainly via physical quenching [lo]. Moreover, we carried out control runs varying the photolysis beam fluence and hence changing initial concentrations of active species. No noticeable changes were observed in the NH2 kinetics, providing additional evidence for the postulated mechanism (5)-(9)The solution of the differential equations corresponding to the scheme (S)-(9) gives for [NH,] in the ground state:
[NH21, = EN&l o exd--k#) + [NH;] 0 [kG/(kG +kc -krIl
where k, = k, [03], X-z = kg [03], k: = kg [Ar] + k7 [NH31 + kg 1031, and INHzl,-,and PHIIo are the initial concentrations of the ground-state and vibrationally excited NH, formed in reaction (2) In accordance with (II) the observed NH, kinetics had a “rise-and-fall” form as shown in fig 2. In the fall region (t > 300 us in our conditions) the observed curves are accurately exponential, giving the reaction (5) pseudo-first-order rate constant k,_ By varying the ozone concentration we have obtained the bimolecular rate constant of the ground state NH2 reaction z;h ozone: kg = (1.5 f 0.3) X lo-l3 cm3 molecule-l The much faster increasing portion of the curve corresponds to the NW2 decay processes (6)-(9). Using the known k, value, we may multiply the eq. (II) by exp(k,t), and so reduce it to a simple saturation curve:
0
0
I 0
8
*
1
1
400
Fig. 2 Ground-state NH2 kinetics Hereafter NH:
denotes viirationally excited
NH2
radi&
Icxp(--krt)
- exp [-(k: + k:) tl1,
Other active species, such as OH(u= 0, 1, 2, 3), 0(3P), O,(lA), were formed in our system along with NH; and NH,. However, it may be shown that they had negligible effect on the NH, kinetics The typical interaction time constants of 0(3P), OH and O&A) with NH2 and NH3, caculated from the reported rate constants and estimated initial concentrations, substantially exceed the measurement timescale. The interaction of vibrationally excited hydroxyl radicals with ammonia also may be neglected because *
1985
[O,]
TIME/&s [NH31
ml
= 9 X 1014 cmm3,
= 1 x 1016 cm-3,Pt&=6Torr.
47
CHEMICAL PHYSICS LEl-l-ERS
Volume 120, number 1
C(r)=C,-,+aCo[ l-
exp(-k&t)]
where C(f) - [NH21t exp(k,t) mental data, c= [I-3210>
co’= mH;lo,
,
(III)
is corrected
experi-
kEm=k;+k:-kFc,,
a=.k*g/k;KEq_ (III) may easily be transformed using semilog coordinates: In [ I - C(r)/(C,
= I,,
[ac; /(co
to a linear form
+ oCg*)] + c&,*)1 -
x-$-rr-
(W
It is to be noted that Co + crc*u represents a maximum level of the ground-state NH, concentration after the relaxation is over, and it may be determined from the corrected experimental curves (III) as the saturation level. The experimental data were plotted using semilogarithmic coordinates [In [I - C(t)/(Co + aCG)] , t] and the relaxation pseudo-first-order rate constant k& was obtained as a slope for each profile. As shown in fig. 3, k&was found to be linearly dependent on [03] through the whole probed [Ox] range, the upper and lower lines corresponding to 1 X 1015 and 5 X 1Or4 crne3 in NH, concentrationS respectively. Thus, the slopes of the lines in fig. 3 give, after correction fork,, the total second-order rate constant of the vibration&y excited NH, radicals removal by
ozone: kg + k8 = (1.5 f 0.4) X lo-r2 cm3 molecule-t s-1. From the difference between the intercepts in fig. 3, we estimate the rate constant of NH; quenching byNH3: k,=9X lo-l2 cm3 molecule-l s-r_ The good agreement of this k, value with Sarkisov et al. [ 1 l] probably-indicates that in our system, as well as in the ammonia photolysis system, the reaction (7) rate is limited by quenching of the low-lying bending mode levels of NH, radicals The rate constant of NH; removal by ozone is too high to be explained simply in terms of V-V transfer because the energy difference between vibrational levels of NH2 and O3 is large - 400 cm-t. So it seems to be of considerable interest to determine the relative importance of reactive and quenching channels. Therefore we have made an attempt to measure the k8/kg ratio. This ratio can be determined here because the relative increase in ground-state NH2 concentration due to the physical quenching depends on the percentage of NH; removed by reaction (9). If, for instance, all the NH; radicals took part in reaction (9), then there would be no increase in [NH2]. We can consider this effect quantitatively as follows: According to (IV), transforming the experimental data to linear form gives In [crC6/(Co + crCi)] as the intercept. The relative portion of NH2 concentration caused by NH; decay fl= oC~/(Cn + oG> may be written as l/P = 1 + Co/C;
= 1 +c,/c,
o-i +-2
0
,
,
4
8
[oJ/iO"cmm'~
~
I2
Fig. 3. Pseudo-&st+rder rate m&ant for vibrationally cxcited NH2 decay x-k versus ozone concentration Ptot = 6 Torr.1--[NHs]=l~1015cm-3,2-[NH~]=5X1014 cmm3_
48
+ (C,-,/C;)(k:
- k,)/k;
+(C&,)
~W~-k~)I0~lI(k~LW
+
’[
27 September 1985
+bINH31 +kg &I)(v)
If there is no vibrational acceleration of reaction (5) (i.e. if kg = k,),the value of P_l does not depend on the ozone concentration But if physical quenching is negligible, B_l grows linearly with [O,]. Fig. 4 shows the dependence of p1 on [03] obtained from our experiments. In spite of scatter in the data the increase of P_l with [03] is clearly visible. The slope of a least-squares line in fig. 4 gives the value of kg - kg, which leads to kg = 1.2 X IO-l2 cm3 s-l_ Thus, the reactive channel (9) is the dominant one and contributes about 80% to the NH; removal by ozone. The comparison of this kg value with the Arrhenius
I
CHEMICAL PHYSICS LETTERS
Volume 120; number 1
‘cl Fig 4. Dependence
I 5
fO-‘5-[0,, cm-’ 10
of p-r on ozone concentration
27 September
1985
conclusion that (35 + 15)s of the exothermicity of reaction (2) (that is 15 i 6 kcal/mole) is distributed among the NH, vibrational modes. This value is in disagreement with the present work. In facf assuming linearity of the suprisal plot (as ’ was done by Kinsey et al.), we estimate the mean vibrational energy of NHi to be 6 kcal/mole using the measured value of the ground-state NH, fraction, 42%. In our opinion, this discrepancy may be accounted for by errors in the energy balance method. Another explanation would be a strong bimodality in the NH, vibrational energy disposal_
(see text). References
expression for the rate constant of reaction (5), k5 = 4 X lo-l2 exp(-llOO/~ cm3 s-1 [7], shows that vibrational energy of NH, may be characterized by a 30% efficiency in overcoming the activation barrier. As yet we do not know which particular viirational mode and level correspond to the acceleration of reaction (5). 3. I lViYz nascent vibrational energy distn_butioIl Extrapolation to zero ozone concentration in fig 4 gives the value of 1 + Co/C,* = 1.7 f 0.2 (see expression (V)). Thus only (442 f 8)% of NH, formed in reaction (2) is in the ground vibrational state. This result may be compared with the data obtained for reaction (2) by Kinsey et al. [ 12]_ They determined the relative populations of u = 0 and u = 1 OH vibrational levels from the measured OH(O) and OH(u= 1) rotational energy disposal. In addition to this, they made approximate measurements of the mean translational energy of the products, and the mean rotational energy of NH,. Energy balance Ied the authors to the
[l]
A.A. logansen, 0-M. Sarkisuv and S.G. Chesk&, Khirn. Phyr (1985) 3-372 [in Russian]. I21 I..% Fletcher and D. Hunti. Can J_ Chem. 54 (1976) 1765. [31 J-A. Davidson, H-1. Schiff, G.E. Streit. J.R. hlcAfee, A.L. Schmeltekopf and CJ. Howard, J. Chem. Phyr 67 (1977) 5021. 141 H. Kurasawa and R. Lcsclaux. Chem Phyr Letters 72 (1980) 437. 151 V.P. Bulatov. A.A. Buloyan, S-G. Cheskis M-2. KozIincr, O.M. Sarkisov and A-1. Trostin, Chem. Phys Letters 74 (1980) 288. 161 W. Hack, 0. Horie and H.Gg. Wagner, Bcr. Dunserges Physik_ Chem 85 (1981) 72. 171 V.P. Buktov. AA. Buloyan, S.G. Cheskis, A.A. Iogansen and O.hl. Sarkosov, Khim Phys (1982) 2-513 [in Russian]. [8] R. Patrick and D-M. Goldeq J. Phyr Chem. 88 (1984) 491. [9] R. Lesclaw Rev. Chem Intermediates 5 (1984) 347. [IO] G.P. Glass H. Endo and B.K. Chaturvcdi J- Chem. Phyr 77 (1982) 5450. [ 1 l] S.G. Cheskh O.hl. Sarkosov and S.Ya Umzmsky, Dokl Akad Nauk SSSR 246 (1979) 66 [inRursian]. 1121 J.F. Cordova, CT. Rettner and J-L. Kinscy, J. Chem Phyr 75 (1981) 2742.
49