J . Phys. Chem. 1985,89, 812-815
812
Reaction of Iodlne wlth Ozone in the Gas Phaset A. C. Vikis* and R. MacFarlane Research Chemistry Branch, Whiteshell Nuclear Research Establishment, Atomic Energy of Canada Limited, Pinawa. Manitoba, Canada ROE 1LO (Received: October 15, 1984)
Elemental iodine (I2) reacts with ozone (0,)in the gas phase to form a solid iodine oxide with the stoichiometriccomposition 1409.The reaction was studied in a flow system, with a N2/02mixture as carrier gas, at a total pressure of 100 kPa and in the temperature range 293-370 K. It was shown by gas-phase titration that 3.9 & 0.2 molecules of O3were consumed per I, molecule reacted. The reaction rate fitted a rate law that was first order in O3 and in I,. The rate constant, k = (-d[Iz]/dt)/([12] [O,]), in the range 293-370 K and in units of dm3mol%-’, was given by In k = (14.7 f 0.6) - (2050 h 230)~’.
1. Introduction The reaction of I, with O3 in the gas phase, or in carbon tetrachloride, has been used in the past for the preparation of iodine oxides.’V2 More recently, Hamilton et aL3 showed that I2 was effective in suppressing the formation of O3 in polluted atmospheres. Also, Glasgow and Willard, in their study of reactions of electronically excited 12,4observed that ground-state I2 reacts consuming -3.7 molecules of 03,per I2molecule. Apart with 03, from the above, there is no information in the literature on the kinetics of the reaction or on the products formed. W e are interested in this reaction as part of a general study to identify and characterize reactions of radioactive iodines (1311 and 1291)with air in nuclear facilities, and in particular because of the potential application of this reaction to remove radioactive iodine from air. Gas-phase conversion of radioactive iodine to solid iodine oxides, which can be removed by particulate filtration, would have several advantages compared to existing removal methods.s 2. Experimental Section The reaction was studied at a total pressure of 100 kPa in the flow system shown in Figure 1. Nitrogen and O2were used as carrier gases for the I,. Oxygen was used as a carrier gas for the 03.The concentration of I2 in the reaction vessel was controlled in the range 106-10-5 mol.dm” by varying the temperature of the fine I, saturator (S,) at fixed I2 carrier-gas flow rates or by varying the ratio of the 1, and O3 carrier gases at a fixed I, saturator (S,) temperature. The I2 was condensed downstream of the reaction vessel in a trap (D), which was packed with 4-mm glass beads and cooled to 195 K with a dry ice/acetone bath. The condensed I2 was dissolved in C C 4 and analyzed spectrophotometrically by using its known absorption at 520 nm. The ozone was generated by passing 0,through a corona mixture entered the reaction vessel discharge (A). The 03/02 (B) through a coaxial 2-mm-i.d. nozzle. The concentration of O3 in the reaction vessel was varied by varying the power to the corona discharge. The O3was monitored immediately past the reaction vessel by light absorption at 253.7 nm. The 253.7-nm radiation was isolated from the spectrum of a low-pressure mercury lamp by a narrow-band interference filter. The reaction vessel was cylindrical (2.2-cm i.d. X 5 5 cm) and made of Pyrex glass. It was thermostated with electrical heating tape wound uniformly around the reaction vessel and sandwiched between two layers of fiberglass insulation. The two carrier-gas streams were also preheated to the temperature of the reaction vessel by similar means. The fiberglass insulation around the reaction vessel also shielded the reacting gases from room light, which was observed to enhance the rate of reaction. This enhancement was assumed to be due to photochemical formation of iodine atoms, followed by reaction with 0,. To minimize losses of Iz and/or O3 by processes other than reaction, Pyrex glass tubing was used throughout the monitored Issued as AECL-8431.
0022-3654/85/2089-0812$01 S O / O
flow system path, except for three greaseless O-ring seals, which connected the reaction vessel outlet and two inlets, and the Teflon plugs of two 4-mm stopcocks used to divert the flow through the I2 trap. Reagent mixing was assumed to be instantaneous in this system. Calculations based on laminar flow showed that, through diffusion alone, the reagents achieve essentially flat radial and axial concentration profiles within the first 2 cm from the ozone inlet. Since the flow is in the turbulent regime (Re lo4), mixing is expected to occur faster than through diffusion in laminar flow. Thermogravimetric (TGA), differential thermal (DTA), and elemental analyses were used for the identification of the solid iodine oxide product. For these purposes the product was scraped off the walls of the reaction vessel and placed in glass vials in a desiccator for subsequent analysis. No visible changes in the samples were observed over several days, provided they were kept in a dry atmosphere. When exposed to ambient air, the samples turned from pale yellow to brown within a few hours. Error values quoted in this text are average deviations for section 3.1 and standard deviations ( f l u ) for sections 3.2 and 3.3.
-
3. Results and Discussion 1. Product Analysis. The TGA and DTA results are shown in Figure 2. Two transitions were observed by using TGA. The first transition at 370-420 K corresponds to a weight loss of 8.1 f 0.3%. More than 98% of the sample’s weight is lost in the second transition at 570-730 K. The DTA data showed an endothermic transition in the same temperature range as the second TGA transition. Due to instrument drift, no DTA transition could be observed in the temperature range of the first TGA transition. However, a sharp exothermic DTA transition was observed at 480 K, which does not appear to have a corresponding TGA transition. The TGA and DTA results obtained in the range 570-730 K agree with those found in previous studies6 on the decomposition of 1205via the following reaction: 21205(s)
-
2 M g ) + 5Odg)
(1)
Wikjord et al.7 performed TGA and DTA on 1 4 0 9 and 1204 formed by passing CH31and air through a negative corona discharge. Our results are in general agreement with their 1 4 0 9 results, except for a small exothermic transition in their DTA curve at 415 K and a shoulder in their TGA curve at 570-730 K, which we did not observe. The former feature could have been missed (1) Brauer, G., Ed. “Handbook of Preparative Inorganic Chemistry”, 2nd ed.; Academic Press: New York, 1963; Vol. I. (2) Emeltius, H. J., Sharp, A. G., Eds. “Advancesin Inorganic Chemistry and Radiochemistry”; Academic Press: New York, 1963; Vol. 5. (3) Hamilton, W. F.; Levine, M.; Simon, E. Science 1963, 140, 190. (4) Glasgow, L. C.; Willard, J. E. J . Phys. Chem. 1973, 77, 1585. (5) ‘Radioiodine Removal in Nuclear Facilities, Methods and Techniques for Normal and Emergency Situations”, IAEA Technical Report Series No. 201, Vienna, 1980. (6) Selte, K.; Kjekshus, A. Acta Chem. Scand. 1968, 22, 3309.
0 1985 American Chemical Society
The Journal of Physical Chemistry, Vol. 89, No. 5, 1985 813
Reaction of Iodine with Ozone in the Gas Phase
= 12
€1
:
?g 10
Figure 1. Reaction flow system: A, ozone generator; SIand S2,coarse and fine iodine saturators; B, reaction vessel; C, ozone detector; D, iodine trap.
3
Figure 3. Ozone consumption vs. iodine flow: 0, the I2 flow was varied by varying the temperature of the fine saturator, S1, at constant carrier-gas flow; 0 , the I2 flow was varied by varying the carrier-gas flow at constant I2 saturator, S2, temperature. TABLE I: Reaction Rate Measurements for the 13/03 Reaction
300
400
500 600 TEMPERATURE / K
700
800
Figure 2. Thermogravimetric and differential thermal analysis curves for the iodine oxide product. Peaks and valleys in the DTA curve correspond to exothermic and endothermic transitions, respectively. X1 and X2 correspond to amplification factors.
by us because of large instrument drift; the latter feature was attributed to particle-size effects.' The origin of the exothermic transition at -480 K, observed in both studies, is uncertain. It could be related to the first TGA transition or to a phase transition of the Iz05,which is formed by the decomposition of the original iodine oxide. Assuming that 1205is produced by the first transition, the TGA results, coupled with the observation that only iodine was released a t the 370-420 K transition (see next paragraph), can be used to calculate the 02:12stoichiometry of the initial oxide, IO,, viz. lOIO,(s) = (2x)I,05(s) + (5 - 2x)I,(g) A value of x = 2.24 f 0.01 was calculated from eq 2 and the 8.1 f 0.3%weight loss a t 370-420 K. The stoichiometry of the initial iodine oxide was also determined by thermal decomposition to its constituent elements (I2 and 0,) in a calibrated evacuated volume (59.7 f 1 cm3). The released iodine was condensed a t 195 K in a cold finger attached to the above volume. No rise in pressure was observed until the sample was heated to about 570 K. The latter observation indicates that only 1, was released in the 370-420 K range. The 0, pressure was measured to within f0.13 Pa with an MKS Baratron vacuum gauge. The I2 was dissolved in CC14 and measured spectrophotometrically by using its absorption at 520 nm. An 02:12stoichiometry of (2.27 f 0.08):l was determined by using this method. The above stoichiometric measurements show that the initial product was 1409. Preliminary studies show that the laser Raman scattering spectrum of the compound is distinct from those of 1,04 and 1205,although there are also several common features. 2. Reaction Stoichiometry. The stoichiometry of the reaction was determined by measuring R , the molar ratio of O3consumed per I2 reacted. These experiments were performed at room tem(7) Wikjord, A. G.; Taylor, p.; Torgerson, D. F.; Hachkowski, L. Thermochim. Acta 1980, 36, 367.
[I2lir [03li, [Ql, pmol. pmolpmol. dm-' dm-3 dm-3 T = 293 K. t = 56.6 s 1.10 18.5 14.4 1.10 15.5 11.7 1.10 14.3 11.6 1.10 11.2 8.59 1.10 11.0 7.71 1.10 8.20 5.96 1.10 5.37 3.51 2.28 17.7 10.1 2.28 13.8 9.02 2.28 11.5 5.63 2.33 9.19 4.40 2.28 8.15 3.28 2.28 5.43 1.52 4.54 15.4 4.24 4.59 13.6 3.58 4.63 13.4 3.05 4.63 11.4 2.14 4.51 10.6 1.61 4.54 9.19 5.30 4.54 7.00 1.69
T = 308 K, t = 53.9 s 1.14 15.6 11.3 1.14 13.2 9.26 1.14 12.5 8.90 1.14 10.1 6.71 1.14 6.76 4.03 2.20 13.1 10.1 2.28 12.7 7.02 2.30 9.62 2.29 2.17 9.47 4.77 2.26 8.74 2.10 2.17 5.67 2.32
tI21i,
[03li>
[OJ,
pmol. pmolpmol. dm-3 dm-) dm-3 T = 323 K. t = 51.4 s 1.12 21.4 17.8 1.12 14.7 10.7 1.09 13.6 9.43 1.09 11.3 7.49 1.12 10.1 6.20 1.09 7.49 4.33 1.09 4.78 2.44 1.09 3.29 1.37 2.09 9.97 3.03 2.11 7.74 1.31 2.17 6.40 0.987 2.01 5.95 0.575 2.07 4.96 0,595 1.99 3.38 0.477 2.03 2.52 0.162 2.05 2.10 0.184 2.07 1.75 0.187
T = 348 K, t = 47.7 s 1.01 4.28 1.61 1.87 11.3 4.13 1.92 10.3 2.94 1.94 5.66 0.608 1.88 4.20 0.432 1.92 1.64 0.0640 1.92 1.36 0.0856 3.90 9.26 0.0566 3.90 6.86 0.0778 3.90 3.45 0.107 T = 370 K, t = 44.8 s 0.884 6.89 4.02 0.884 3.70 0.408 0.884 2.49 0.354 0.884 2.39 0.333 1.81 6.89 0.291 1.76 5.50 0.0862 1.76 4.27 0.133 1.81 3.63 0.0468 1.81 2.83 0.0801 1.81 2.12 0.0134
perature, using oxygen as the carrier gas for both 1, and 03,under conditions of more than 99% I2 removal. The 1, concentration was varied by varying the flow through the fine I, saturator, at constant saturator temperature, and by varying the temperature of the fine 1, saturator at constant flow. The results from the two methods were essentially identical and are shown in Figure 3. A value of 3.9 f 0.2, calculated from the slope in Figure 3, was determined for the 03:1, reaction stoichiometry. 3. Reaction Rate and Temperature Dependence. The reaction for various initial rate data obtained for the consumption of 03, O3 and I, concentrations and temperatures, are summarized in Table I. Due to experimental limitations, only a narrow range
814 The Journal of Physical Chemistry, Vol. 89, No. 5, 1985
Vikis and MacFarlane
IODINE / OZONE
O r 0 , 0 , 2 10,4,
00 0 4 oa
? 0 , 0 1 4 , 0 1 8 l,2 q
-12
\
Or
I
O
I
I\
21 LEGEND 0
-30
-
1
293K
A
T-I
370K
00 0 2 0 4
IODINE/OZONE
Figure 4. Reaction rate vs. [Iz]/[03] ( Y = ([O,]it)-' In [([121i[03],)/ ([I21r[O3li)l).
TABLE II: Rate Constants for the 12/03 Reaction temv. K k f u. dm3.mol-'-s-' 293 308 323 348 370
(2.25 f 0.19) X lo3 (3.00 f 0.28) X lo3 (5.72 f 0.57) X 10' (9.85 f 1.56) X IO3 (1.74 f 0.28) X lo4
R f u 4.0 f 0.5 3.9 f 0.6 4.5 f 0.5 3.9 f 0.7 4.3 f 0.8
-d[I,]/dt = -R-'d[O,]/dt = k[I2][0,]
(3)
where k is the reaction rate constant and R is the reaction stoichiometric factor. (It should be noted here that more complex rate laws would also provide a reasonable fit to the rate data; however, as the following analysis shows, the simple rate law chosen has some theoretical justification.) The integrated rate law, from eq 3, is
where the subscripts i and t refer to initial and final (at time t ) concentrations, respectively. To determine the reaction rate constant, the left-hand side of eq 4 was divided by [O,lit and the resulting expression
was plotted against the ratio [12]i/[03]i,as shown in Figure 4. The final iodine concentrations were calculated from the reaction stoichiometry, with R = 3.9, as determined in section 3.2. At each temperature, the rate constant was determined from the intercept of the linear least-squares fit of the data. The slope to intercept ratio is also equal to t h e value ( R ) of t h e reaction stoichiometry. Values of k and R, determined from the linear least-squares fit, are given in Table 11. The latter values of R are (within the error limits) temperature independent and compare well with the room-tempeature value determined in section 3.2. The rate constants were fitted to the Arrhenius equation ( k = Ae-E*/Rq.A preexponential factor ( A ) of 107.76M.17 dm3.mol-1-s-1 and an activation energy (E,) of 25.0 1.2 kJ.mol-I were determined by a linear least-squares fit of the data (see Figure 5 ) , using the logarithmic form of the Arrhenius equation (In k = In A - EJRT). The magnitude of the preexponential factor is of the order of magnitude expected for bimolecular reactions between diatomic and nonlinear triatomic molecules, according to transition-state theory.* A steric factor of 10" is predicted for such reactions,
*
1 1 6 ~ ~
Figure 5. An Arrhenius plot of the rate constants. Vertical error bars correspond to the standard deviation of each rate constant. Horizontal error bars correspond to the average error in the temperature measurement.
and a value of -5 X lo4 can be estimated for the 12/03reaction. There is also a close analogy between this reaction and the reactions of O3with olefins, for which similar preexponential factors were determined.g*'O Thus, a cyclic transition state, as proposed for the 03/olefin reactions? may also be appropriate for the 12/03 reaction. Two feasible reaction channels are
of concentrations at an essentially fixed reaction time could be studied. Thus, the rate law could not be established from measurements of initial rates. Instead, a rate law that was first order in [I2] and in [O,] was assumed
-
L A 2.6 2.8 3.0 3.2 3.4 3.6
4-40
I, t 0,
-
:dy\oj -
-1, IO t I t
1
1-1
1
2
02
(5)
IO t 102
Channel 1 is endothermic by 32 i 13 kJ-mol-', and channel 2 is estimated to be less endothermic. (The endothermicity was calculated by using the IO bond dissociation energy of 222 f 13 kEmol-', as determined by Radlein et al." Earlier literature valuesI2 are lower by about 42 kJ-mol-' and result in an endothermicity of 74 kJ.mol-', which is too high to have permitted a measurable rate via these reactions channels.) The observed activation energy is consistent with these endothermicities. The relative contributions of the two reaction channels cannot be established from the present work. However, it is interesting to is formed by successive observe that if the final product (1409) oxidation of the I, IO, and IO2 primary products, i.e.
I
+ O3
IO
-
+ O3
IO
+ O2
IO2 + 02,etc.
(6)
(7)
then at least 4.5 molecules of O3per I2 molecule would be required if only channel 1 were available. The corresponding stoichiometry, if only channel 2 were available, is 2.5. The above stoichiometries are minimum values, because no account is taken of catalytic destruction of O3 by reactions such as
IO
-+ -+
+ O3
210
I
12
202
0 2
(8) (9)
Thus, in view of the observed stoichiometry, a significant fraction of the reaction probably w u r s via channel 2, which is also more favored energetically. Two alternative reaction mechanisms have also been considered for the 12/03 reaction. These are thermal decomposition of 03, followed by 0 atom reaction with I2 03
+M e 0 2+0+M
(10)
(8) Pratt, G. L. "Gas Kinetics"; Wiley: London, 1969. (9) DeMore, W. B. Znt. J . Cfiem. Kinet. 1969, 1 , 209. (10) Huie, R.E.; Herron, J. T. Znt. J . Chem. Kinet., Symp. 1975, No. 1 . (1 1) Radlein, D. St. A. G.; Whitehead, J. C.; Grice, R. Nature (London)
1975, 253, 37. (12) Darwent, B. deB. Natl. Stand. Ref. Data Ser. (US.Nutl. Bur. Stand.) 1970, NSRDS-NBS 31.
0
+ I2
-
J. Phys. Chem. 1985, 89, 815-820 products
(11)
and thermal decomposition of 12, followed by I atom reaction with 0 3
+M I + O3 I2
-
+M
(12)
products
(13)
e 21
Thermal decomposition of O3is too slow in this temperature range to account for the observed reaction rates. The activation energy for the forward step of reaction 10 is 103 kJ.m01-'.'~ According to the data of Benson and Axworthy,I3 for the reaction times used in our work, about 0.01% of the initial O3would have reacted at 293 K and about 7% at 370 K. Reaction via the thermal decomposition of I2 would be even slower, in the absence of light, because an activation energy equal to or greater than the bond dissociation energy of I2 (1 5 1 kJ-mol-I) would be required. Thus, atomic reaction mechanisms cannot account for the observed rate data. (13) Benson, S.W.; Axworthy, Jr., A. E. Adu. Chem. Ser. 1959, No. 21, 398.
815
4. Conclusions Elemental iodine was shown to react with ozone, forming a solid iodine oxide with the stoichiometric composition 1409. The stoichior;netry, rate, and temperature dependence of this reaction were studied in the range 293-370 K. The observed kinetics suggest that the rate-determining step is a bimolecular reaction between I2 and O3to form I, IO, and IO2 radicals, which undergo further oxidation by reaction with O3 to form the final 1409 product. The reaction rate and product stability are suitable for application of this reaction to the removal of radioactive iodines from air in nuclear facilities. In the presence of ionizing radiation, such as in a nuclear reactor accident, where some ozone would be produced by radiolytic decomposition of the oxygen in the air, the reaction of I2 with O3 is expected to reduce the potential releases of volatile radioiodines. Acknowledgment. We thank Ms. D. A. Furst and Mr. B. D. Wilson for their assistance and Drs. D. F. Torgerson and P. Taylor for helpful discussions. Registry No. 12, 7553-56-2; Oj, 10028-15-6.
Photoinduced Isomerization of Radical Ions. 3. Radical Cations of Cyclopentadiene, Dlcyclopentadlenes, and 1,3-Blshomocubane Produced in 7-1rradlated Freon Matrices at 77 K Tadamasa Shida,* Takamasa Momose, and Noboru Ono Department of Chemistry, Faculty of Science, Kyoto University, Kyoto 606, Japan (Received: December 7, 1984)
The radical cations of cyclopentadiene and the three CI0Hl2systems in the title have been produced in y-irradiated Freon matrices at 77 K. Electronic and ESR spectroscopy were used to determine the interrelationships among the cations and their photoproducts. The three Cl&Il2+.'syielded different photoproducts depending upon the wavelength used for photoexcitation. The photoproduct obtained from endo-dicyclopentadienecation excited to its first excited state was also obtained by a thermal reaction between the neutral molecule of cyclopentadiene and its radical cation. The product is inferred to be the radical isomer are discussed. cation of a new dicyclopentadiene. Possible mechanisms of obtaining this new CloHI2+*
1. Introduction y-Irradiation of frozen solutions using Freons as the solvent leads to the production of solute radical cations.' A glass-forming admixture of CC13F and CF2BrCF2Br (1: 1 by volume) is used for the optical studies because of its transparency in the near-UV and near-IR regions.24 For the ESR study a polycrystalline matrix of CC13Fis found useful.',5 In this serial work we present the result of studies on the radical cations in the title as well as that of tram-tricyclo [3.0.3 .O]deca-2,8-diene. They were produced by irradiation in the Freon mixture mentioned above (abbreviated as F M hereafter). In a separate paper we have discussed the nature of the electronic absorption spectra of the radical cations of end+ and exo-dicyclopentadienes in FM.6 In the present paper (1) Shida, T.; Haselbach, E.; Bally, T. Acc. Chem. Res. 1984, 17, 180. (2) Sandorfy, C. Can. J. Spectrosc. 1965, 10, 85. (3) Grimison, A.; Simpson, G. A. J . Phys. Chem. 1968, 72, 1776. (4) (a) Shida, T.; Iwata, S. J. Am. Chem. SOC.1973,95, 3473. (b) Shida, T.; Kato, T.; N m k a , Y . J. Phys. Chem. 1977,81, 1095. (c) Shida, T. J. Phys. Chem. 1977,82,991. b and c constitute the first and second parts of this serial work. (d) Shida, T.; Nosaka, Y.;Kato, T. J. Phys. Chem. 1978,82,695. (e) Kato, T.; Shida, T. J. Am. Chem. SOC.1979, 101, 6869. ( 5 ) Shida, T.; Kato, T. Chem. Phys. Lett. 1979, 68,106. (6) Momose, T.; Shida, T.; Kobayashi, T., to be submitted for publication.
our focus will be mainly on the spectral change induced upon photoexcitation. The dependence of the optical spectrum upon the initial concentration of cyclopentadiene was also examined. Subsidiary ESR studies were carried out briefly. Complete identification of all the photoproducts was not achieved owing to the complexity of the system. However, the photoinduced reactions are consistently interrelated and plausible reaction mechanisms are considered (see Scheme I). 2. Experimental Section endo-Dicyclopentadiene was commercially obtained. Cyclopentadiene was produced therefrom by the standard method. em-Dicyclopentadiene and 1,3-bishomocubane were synthesized and isolated by literature procedures.'~* A less stable dicyclowas also synpentadiene, truns-tricyclo[3.0.3.0]deca-2,8-diene, thesized as an admixture with endo- and ~ X O - D C P . ~ ~ ' ~ (7) Nelson, G. L.; Kuo, C.-L. Synthesis 1975, 105. Steinmetz, R. Chem. Ber. 1963, 96, 520. (8) Schenck, G. 0.; (9) Turro, N. J.; Hammond, G. S. J. Am. Chem. SOC.1962, 84, 2841. (10) Hammond, G. S.; Turro, N. J.; Liu, R. S. H. J . Org. Chem. 1963, 28, 3297.
0022-3654/85/2089-0815$01.50/00 1985 American Chemical Society
