The Journal of
Physical Chemistry
0 Copyright, 1982, by the American Chemical Society
VOLUME 86, NUMBER 24
NOVEMBER 25, 1982
LETTERS Chlorine Perchlorate. A Major Photolysis Product of Chlorine Dioxide A. J. Schell-Sorokln, D. S. BethuWt J. R. Lankard,+M. M. 1.Loy,z and P. P. Sorokln't IBM Thomas J. Watson Research Center, YorMown Heights, New York 10598 (Received: June 8, 1982; I n Final Form: September 7, 1982)
Gaseous chlorine perchlorate, C10C103, is found to be a major photolysis product of chlorine dioxide (OClO). Experiments were performed at room temperature with both continuous wave (mercury lamp) and pulsed (XeCl UV laser) light sources. The technique of time-resolved IR spectral photography (TRISP) was used to monitor the growth of the strong 1282-cm-l C10C103IR band following the application of a single -35 ns long, -70 mJ/cm2, XeCl laser pulse to a mixture of -30 torr of OClO and -700 torr of N2 It was found that under these conditions this band forms with a time constant of -1 ps. A transient IR band of unknown origin at N 1232 cm-I was also observed to develop on the same time scale.
Several photolysis studies of chlorine dioxide (OClO) have been conducted since its discovery in 1815 by Sir Humphrey Davy.l By photolysis of chlorine dioxide, Millon2 in 1843 first synthesized a larger chlorine oxide, C1206. This red oily liquid, chlorine hexoxide, was rediscovered in 1925 in chlorine dioxide photolysis studies by Booth and Bowen3 and by Bodenstein, Harteck, and Padelt.* In 1967 McHale and von Elbes reported that a brownish substance condensed when chlorine dioxide gas at -45 "C was irradiated by light from a mercury lamp. The brownish condensate was shown to be a mixture of C1206and an unknown, more volatile, substance. Warming the condensate sliahtlv under vacuum caused the more volatile fraction to gasify and decompose instantly to chlorine and oxygen in a 2:3 ratio. This finding led the authors of ref 5 to believe that they had isolated a new oxide of chlorine, chlorine sesquioxide, C1203. Further'Supported in part by the US.Army Research Office. 3 Supported in part by the Office of Naval Research. 0022-3654/82/2086-4653$01.25/0
more, by evaporating small amounts of the new substance into chlorine dioxide vapor they observed drastic reductions in explosion induction times when the mixtures were subsequently admitted to heated bulbs! McHale and von Elbe6 thus concluded that Cl,03 was the intermediate responsible for the delayed explosive reaction, and they formulated a branched chain kinetic model based upon this assumption. However, no spectral data of any kind were observed in these experiments, nor have corroborative studies of these effects been published since. We report a new product, also a large chlorine oxide, formed in the room temperature photolysis of chlorine (1)H. Davy, Phil. Trans., 105, 214-9 (1815). (2)E. Millon, Annales, 45, 281;46, 312 (1843). (3)H. booth and E. J. Bowen, J. Chem. SOC.,127,510-3 (1925). (4)M.Bodenstein, P.Harteck, and E. Padelt, Z. Anorg. Allg. Chem., 147, 233-44 (1925). (5)E. T.McHale and G. von Elbe, J. Am. Chem. SOC.,89, 2795-7 (1967). (6)E.T.McHale and G. von Elbe, J. Phys. Chem., 72,1849-56 (1968).
0 1982 American Chemical Society
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The Journal of Physical Chemlstty, Vol. 86, No. 24, 7982
dioxide. The new photolysis product is gaseous chlorine perchlorate, C10C1O3 This compound was discovered and first synthesized in 1970 by Schack and Pilipovich.7 These authors included in their description of chlorine perchlorate its infrared spectrum which permitted its identification as the photolysis product observed here. We found that chlorine perchlorate is produced when 0.25-40 ~OITof chlorine dioxide is mixed with a buffer gas (nitrogen or C02) to a total pressure of 760 torr and irradiated in a 700-cm3stainless steel reactor with light from a mercury arc lamp. This is the first reported synthesis by a photolytic method of chlorine perchlorate. With use of a filter (Corning 3-75) which transmits wavelengths longer than 390 nm approximately two thirds of the reacted chlorine dioxide forms chlorine perchlorate (as determined by changes in the optical densities of the infrared absorption bands7"), even at the lowest chlorine dioxide pressures utilized (-0.25 torr). With a 10oO-W lamp all the chlorine dioxide (initial pressure -30 torr) is consumed within 20 min. Large amounts of chlorine perchlorate are again formed in the stainleas steel reactor with the use of various ultraviolet pass band filters (Corning 7-51, 7-39, or 7-37, transmitting the ranges 310-410,320-400, and 330-380 nm, respectively). Finally, when the gas is irradiated without a buffer gas, chlorine perchlorate is again formed, but its concentration soon reaches a small steady-state value. Chlorine perchlorate is the only product we observe in the infrared spectrum of a sample of chlorine dioxide photolyzed in a stainless steel vessel at room temperature with the mercury arc lamp, if care is taken to avoid the presence of an unknown H-containing species, probably HC1 or water. The presence of such species leads to formation of HOC103 by reaction with the C10C103.719 In view of the above finding, we were surprised to note the complete absence of chlorine perchlorate or, indeed, of any gaseous product possessing an IR spectrum when the mercury lamp photolysis described above was conducted in a pyrex vessel equipped with NaCl windows. Thus it is perhaps possible to understand why C10C103 has never been seen before in OClO photolysis. A glass or p y e x surface evidently strongly catalyzes the destruction of this particular compound. According to ref 7, chlorine hexoxide is a major decomposition product of chlorine perchlorate. Again, with reference to earlier work, chlorine perchlorate is known to have a vapor pressure -8 torr at -45 "C and, since it sublimes readily at that temperature, it most likely would have been pumped away by McHale and von Elbe5a when they evacuated their reaction vessel prior to evaporation of the brown condensate. The photolytic reaction producing CIOCIOBfrom OClO also occurs at very high UV light intensities. We observed the formation of the same end products (C10C103 and HOC103),again in the stainless steel reactor, when gaseous mixtures of chlorine dioxide (-10 torr) and nitrogen (750 torr) were irradiated with multiple pulses of UV light from an XeCl excimer laser (A = 308 nm). Each laser pulse was -35 ne in duration. Per pulse UV excitation fluences of up to 70 mJ/cm2 were applied. To try to obtain a time-resolved picture of the formation of C10C103,we utilized the recently developed technique of time-resolved infrared spectral photography (TRISP).'o~ll In the 7-10-pm range, the time resolution (7) C. J. Schack and D. Philipovich, Znorg. Chem., 9, 1387-90 (1970). (8) A. H. Nielsen and P. J. H. Woltz, J. Chem. Phys., 20, 1878-83 (1952). (9) K. 0. Christe, C. J. Schack, and E. C. Curtis, Inorg. Chem., 10, 1589-93 (1971). (10) Ph. Avouris, D. S. Bethune, J. R. Lankard, J. A. Ors, and P. P. Sorokin, J. Chem. Phys., 74, 2304-12 (1981).
Letters
of this technique is currently about 1 ns.12 We observed the appearance of the 1282-cm-' absorption band of chlorine perchlorate within 1ps of a 35 nsec, 70 mJ/cm2, XeCl photolyzing pulse applied to 30 torr of chlorine dioxide in an atmosphere of N2 A W quartz cell with BaF2 windows was used. On the same time scale, a second, equally intense, but as yet unidentified band at 1232 cm-' develops. During the time interval 5-50 ps the 1232-cm-' band broadens considerably and appears to weaken somewhat, while the 1282-cm-' C10C103band further intensifies. From -50 ps to -5 ms there is no further change in the appearance of these two bands. At times somewhat greater than -5 ms, under the particular irradiation conditions described here, the gaseous mixture explodes, and both the 1282- and 1232-cm-' bands disappear. The species absorbing at 1232 cm-' is never seen as a final product, even under conditions where no explosion occurs. Finally, we have determined that the strength of the 1232-cm-' band at its maximum becomes intensified by more than a factor of 2 when 160 torr of C12 is added to the OC10-N2 mixture. Correspondingly, the strength of the 1282-cm-' band becomes weaker than in the case with no added Clz gas. At first glance, this might lead one to believe that the 1232-cm-' band is induced by the presence of C1 atoms. However, the reaction
c1+ OClO -% 2c10 is to occur extremely rapidly ( K , = 5.9 X lo-" cm3 molecule-' s-'1. Thus all C1 atoms formed should be converted to C10 radicals in -16 ns under the present circumstances. Hence, it would seem to follow that the 1232-cm-' band appears when C10 radicals are created in the presence of OClO molecules, with a large amount of buffer gas also being present. Besides the two bands at 1282 and 1232 cm-', no other transient IR bands were seen in the 7-10-p range, the only range examined with TRISP in this study. A transient broadband UV absorption spectrum that may correspond to one or both of the species whose IR bands are discussed above was also seen. Specifically, a 3 cm X 2 cm X 1cm quartz cell was filled with a 10 torr of OC10, -750 torr of Nz mixture and was then subjected to a single -70 mJ/cm2 XeCl laser pulse across the entire 3 cm X 2 cm face. Broadband probe light from a pulsed (-3 pa duration) Xe flashlamp was then passed along the long dimension of the cell, dispersed in a spectrograph, and photographed. With the limited time resolution of this optical arrangement three distinct absorption systems were observed: (a) C10 A 2 n X211bands, (b) O2SchumannRunge bands originating from vibrational levels as highly excited as u" = 13, and finally, (c) a new featureless UV 290 nm and exabsorption spectrum originating at X tending to beyond -230 nm. The C10 absorption largely decayed in -50 ~ s The . vibrationally excited O2 disappeared well before this time. The third spectrum which had already become apparent -5 pa after the excimer laser pulse remained basically unchanged in character, although weakened, by -2 ms, the longest probe delay tried. Our observations do not provide enough constraints to allow the choice between several plausible mechanisms for producing both the C10C103and the species absorbing at 1232 cm-l, even when known rate constants are utilized.
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(11) D. S. Bethune, J. R. Lankard, P. P. Sorokin, A. J. Schell-Sorokin, R. M. Plecenik, and Ph. Avouris, J . Chem. Phys., 7 5 , 2231-6 (1981). (12) D. S. Bethune, J. R. Lankard, and P. P. Sorokin, Paper FA5 in the 1981 Annual Meeting Program of the Optical Society of America. (13) P. P. Bemand, M. A. A. Clyne, and R. T. Watson, J.Chem. So;., Faraday Trans. 1,69, 1356-74 (1973).
The Journal of Physical Chemistry, Vol. 86, No. 24, 1982 4655
Letters
To illustrate, let us start by assuming tentatively that a C10C103 molecule results from a bimolecular collision between a ground-state OClO molecule and an energetic precursor molecule (X) having the same chemical formula as OC10. Some possible mechanisms for producing X arise from the assumption that OClO predissociates, as several spectroscopic studies14-16have concluded from the-diffusepess of the lines making up the ultraviolet OC10(A2A2 X2BJ absorption band. The effective dissociation channel for OClO is still totally unknown, however. One possibility is that the fragments are C1 + 02*,with 02* being either the 'Zg+ or 'Ag excited electronic states, or a vibrationally excited level of the ground state 32;. Since with 308-nm photolysis -90 kcal of energy would be could serve as an energy available for the fragments, 02* carrier, creating the precursor X by exciting one of the low-lying excited electronic states (2Al,2B2) of OClO that are predictedI7 to exist: 02*+ OC10(2B1) O2 + OC10*(2A1, 2B2). By demonstrating that C10C103 is formed photolytically with high efficiency in the presence of an atmosphere of C02 buffer gas, we have effectively ruled out the possibility that 02('Zg+)is the carrier, since COz is known to be an effective quencher of this state.'* The possibility that 02('A ) or vibrationally excited 31;; is the carrier remains. bimilarly, if the dissociation products of photoexcited OClO are C10 + 0, as is more generally believed, 02*would again be implicated as an energy carrier. The rate constants for the two reactions by which it would be formed in this case are known.13
-
-
0 + OClO -% 0 2 * k2 = 5
X
+ c10
cm3 molecule-' s-'
0 + c10 -% 02*
+ c1
Iz3 = 5 x lo-'' cm3 molecule-'s-l
The former reaction would be effective in both continuous wave and laser photolysis experiments; the latter reaction would be important only in laser photolysis at high fluence (14)W.Finkelnburg and H. J. Schumacher, 2.Phys. Chem. Bodenstein-Festbond,704-16 (1931). (15)P.A. McDonald and K. K. Innes, Chem. Phys. Lett., 59,5624 (1978). ' (16)S.Michielsen, A. J. Merer, S. A. Rice, F. A. Novak, K. F. Freed, and Y. Hamada, J. Chem. Phys., 74, 3084-101 (1981). (17)J. L.Gole, J. Phys. Chem., 84, 1333-40 (1980). (18)R.G.AvilBs, D. F. Muller, and P. L. Houston, Appl. Phys. Lett., 37, 35840 (1980).
levels (-70 mJ/cm2). The rates calculated from k1,k2are compatible with the TRISP results. A plausible mechanism for producing the 1232-cm-' absorber can also be given. One starts by noting that all C1 atoms formed either directly or indirectly in the two OClO dissociation channels discussed above will be rapidly converted to C10 radicals. This was discussed earlier. An interesting speculation now arises that C10 radicals can, in turn, react with OClO molecules to form chlorine sesquioxide-the compound of McHale and von Elbe5*6and that this, in fact, is the identity of the 1232-cm-' absorber! There is the added possibility that chlorine sesquioxide produced in this manner reacts further with the parent compound to produce chlorine perchlorate in a chain sustaining reaction: cl2o3 O C ~ O ~ 1 0 ~ 1 c0 i ~o
+
-
+
This may account for the changes in the two IR bands observed during the time interval 5-50 ps. The gas-phase reaction producing C10C103that we report here can evidently also occur as a solid-state photolytic reaction. In their matrix isolation study of the ClOO radical made 3 years before the first report of chlorine perchlorate, Arkell and Schwagerlg observed a set of infrared absorption bands and assigned them to unidentified photolysis products of OClO aggregates because the bands had a larger optical density in samples with a larger ratio of OClO to argon. The infrared frequencies and relative band intensities reported by them, in fact, correspond to those obtained by Christe, Schack, and Curtisgfor matrix isolated chlorine perchlorate. The frequenbies in cm-' and relative optical densities of the group of unidentified bands observed by Arkell and Schwager are 564 (0.0231, 586 (O.Oll), 638 (0.048), 648 (0.160), 1037 (0.076) (attributed by the authors to be due solely to ozone), 1258 (0.010),1279 (0.090), and 1290 (0.103). Those of matrix isolated chlorine perchlorate reported by Christe, Schack, and Curtis are 561 (ms), 582 (m), 638.5 (vs), 647 (vs), 1039 (s), 1256 (m), 1271 (vs), and 1287 (vs). The only bands of medium or greater absorption strength of chlorine perchlorate in the spectral region covered in the study of Arkell and Schwager and which are not listed by them are 746 (m) and 513 (m). Qualitative agreement is evident. We are continuing to study this interesting system and the possible role the transients we observe may have in laser-induced explosions of chlorine dioxide. (19)A. Arkell and I. Schwager, J. Am. Chem. Soc., 89,5999-6006 (1967).
