15130
J. Phys. Chem. 1996, 100, 15130-15136
Rate Constants for the Reactions I + OClO, I + ClO, Cl + I2, and Cl + IO and Heat of Formation of IO Radicals Yuri Bedjanian,† Georges Le Bras, and Gilles Poulet* Laboratoire de Combustion et Syste` mes Re´ actifs, CNRS and UniVersite´ d’Orle´ ans, 45071 Orle´ ans, Cedex 2, France ReceiVed: March 6, 1996; In Final Form: June 20, 1996X
The kinetics of the reactions of iodine atoms with OClO and ClO and of chlorine atoms with I2 and IO have been studied using the discharge-flow mass spectrometric method at a total pressure of about 1 Torr. The rate constant obtained for the reaction I + OClO f IO + ClO (2) in the temperature range (288-353) K is k2 ) (9.7 ( 3.0) × 10-12 exp(-(1190 ( 200)/T) cm3 molecule-1 s-1. IO and ClO radicals have been detected as reaction products, and a branching ratio of unity has been found for the channel forming these two radicals. For the reaction I + ClO f products (3), only upper limits of the rate constant have been determined: k3 e 3.7 × 10-15 and k3 e 1.7 × 10-14 cm3 molecule-1 s-1 at T ) 297 and 353 K, respectively. For the reaction between Cl atoms and IO radicals, Cl + IO f I + ClO (4), the rate constant has been observed to be independent of temperature in the range 290-355 K: k4 ) (4.4 ( 1.0) × 10-11 cm3 molecule-1 s-1. Finally, the rate constant obtained for the reaction Cl + I2 f I + ICl (5) is k5 ) (2.1 ( 0.3) × 10-10 cm3 molecule-1 s-1 in the temperature range 296-365 K. From these kinetic data, upper and lower limits for the enthalpy of IO formation at T ) 298 K could be derived: 25.8 kcal mol-1 e ∆Hf(IO) e 27.5 kcal mol-1.
Introduction Interest in the chemistry of iodine-containing species in the atmosphere was initially related to the possible role of iodine in marine boundary layer chemistry.1-5 The potential stratospheric impact of iodine-containing compounds has not been considered until recently, since these compounds are known to be too short-lived to be transported to the stratosphere. For example, CH3I, the major iodocarbon with an oceanic source has a lifetime, calculated from its photochemical loss in the troposphere, of less than 4 days. However, recent views of atmospheric dynamics indicate that transport through deep convection, particularly in tropical regions, could be efficient to transport such short-lived species from the surface to the troposphere.6 It has been calculated that even very small amounts of iodine species reaching the stratosphere (compared to Cl- and Br-containing species) could have a significant influence on stratospheric ozone, especially in the lower stratosphere. One of the catalytical cycles potentially important for ozone depletion, which has been considered in the model calculation of ref 6 involves the interhalogen reaction of IO with ClO. This reaction may proceed through several channels:
IO + ClO f I + OClO ∆H ) -3.0 ( 1.9 kcal mol-1 f I + ClOO ∆H ) -2.3 ( 1.9 kcal mol-1 f ICl + O2 ∆H ) -46.5 ( 0.9 kcal mol-1 f Cl + OIO ∆H ) ?
(1a) (1b)
(1c) (1d)
The values given for the reaction enthalpies are derived from the enthalpy data of the literature,7 except for IO radicals, for # Permanent address: Institute of Chemical Physics, National Academy of Sciences, 5/2 P. Sevak, 375044 Yerevan, Armenia X Abstract published in AdVance ACS Abstracts, August 15, 1996.
S0022-3654(96)00696-X CCC: $12.00
which the value has been obtained from the present work: ∆Hf(IO) ) 26.7 ( 0.9 kcal mol-1. Preliminary kinetic data for reaction 1 have been reported very recently.8 The rate constant of this reaction has been measured as a function of temperature between 200 and 360 K: k1 ) (6.1 ( 1.6) × 10-12 exp[(280 ( 60)/T] cm3 molecule-1 s-1. However, no mechanistic information has been published so far for reaction 1. In the present work, the kinetic parameters for several reactions between iodine and chlorine species have been measured:
I + OClO f IO + ClO ∆H ) 3.0 ( 1.9 kcal mol-1 I + ClO f products Cl + IO f I + ClO ∆H ) -5.7 ( 0.9 kcal mol-1 Cl + I2 f I + ICl ∆H ) -13.7 kcal mol-1
(2) (3) (4) (5)
These measurements were necessary not only to improve the kinetic data base, but also because they were needed prior to the laboratory investigation of reaction 1 since the above reactions are anticipated to be secondary reactions in the chemical systems used to study reaction 1. For example, such systems may use simultaneously I2 (as the precursor of IO, from the reaction of I2 with O atoms) and OClO (as the precursor of ClO, from the reactions of OClO with Cl or O atoms). The other aim of the present study was to derive, from the kinetic data obtained, the enthalpy of formation of IO radicals, which is poorly known so far. Indeed, reaction 2 is the reverse reaction of reaction 1a, which is a possible channel of reaction 1. Thus, the determination of both k2 and k1a should allow for the calculation of ∆Hf(IO). However, this cannot be made precisely since k1a has not been measured directly. Nevertheless, the determination of k2 combined with the recently measured overall rate constant of reaction 1,8 which represents an upper limit of k1a, will give an upper limit for ∆Hf(IO). Similarly, the upper limit measured for k3, which is also an upper limit © 1996 American Chemical Society
Kinetics of Reactions of I and of Cl Atoms
J. Phys. Chem., Vol. 100, No. 37, 1996 15131 titration reaction with Br2 ([Cl] ) [Br2]) gave the same results (within 7%):
Cl + Br2 f Br + BrCl -10
k7 ) 2.4 × 10
(7) -1 -1 11
3
exp(-144/T) cm molecule
s
OClO was synthesized in the laboratory and diluted with He and was kept at a temperature of about 280 K in a flask maintained in the dark. The purity of OClO was determined mass spectrometrically by monitoring both OClO and its main decomposition product, Cl2. ClO radicals were generated by the reaction
Cl + OClO f ClO + ClO (8) k8 ) 3.4 × 10-11 exp(160/T) cm3 molecule-1 s-1 7
Figure 1. Diagrams of the apparatus used: (a) for the study of the reactions of I with OClO (reaction 2) and ClO (reaction 3) and (b) for the study of the reactions of Cl with IO (reaction 4) and I2 (reaction 5).
for the rate constant of the channel forming Cl and IO, combined with the determination of the rate constant for the reverse reaction 4, also allows for the determination of a lower limit for ∆Hf(IO). Experimental Section Experiments were carried out using a molecular beam mass spectrometer coupled to a discharge-flow system which has been described earlier.9 The reactor consisted of a Pyrex tube (45 cm length and 2.4 cm i.d.) with a jacket for the thermostated liquid circulation. Two different configurations of the movable double injector for the introduction of the reactants into the reactor (Figure 1) were used in this study. To reduce the wall loss of active species, the inner surfaces of the reactor and of the two tubes of the movable central injector were coated with halocarbon wax. Iodine atoms were produced from the reaction of molecular iodine with either chlorine or hydrogen atoms:
Cl + I2 f I + ICl
(5)
k5 ) (2.1 ( 0.3) × 10-10 cm3 molecule-1 s-1 H + I2 f I + HI
(6)
k6 ) 6.6 × 10-10 exp(-20/T) cm3 molecule-1 s-1 10 The value of k5 has been determined in this work (see below). H and Cl atoms were produced in a microwave discharge of H2/He and Cl2/He mixtures, respectively. The direct dissociation of I2 molecules in the microwave discharge to generate I atoms was not used since the formation of a deposit at the inner surface of the discharge tube was observed leading to instability in the production of I atoms. I2 was introduced into the reactor by flowing helium through a column containing iodine crystals. The absolute concentrations of both atomic and molecular iodine were derived from the titration of Cl by an excess of I2 (reaction 5): [I] ) [I2] ) [Cl]0 (initial Cl concentration). The mass spectrometric signal for Cl atoms was calibrated by the measurement of the dissociated fraction of known Cl2 concentration passing through the microwave discharge ([Cl] ) 2[Cl2]). It was verified that the absolute calibration of Cl using the
This reaction was also used to determine the absolute concentration of ClO radicals by measuring the OClO concentration consumed by Cl atoms in excess. IO radicals were produced by two different methods using the reaction of oxygen atoms with either molecular iodine or trifluoromethyl iodide:
O + I2 f IO + I -10
k9 ) 1.4 × 10
(9) -1 -1 7
exp((0 ( 250)/T) cm molecule 3
s
O + CF3I f IO + CF3 -11
k10 ) 1.16 × 10
3
(10) -1 -1 12
cm molecule
s
The concentrations of IO radicals were determined using their titration by NO and the absolute mass spectrometric calibration of NO2 formed in reaction 11:
IO + NO f I + NO2 -12
k11 ) 7.3 × 10
(11) -1 -1 7
exp((330 ( 150)/T) cm molecule 3
s
All the calibration experiments were carried out under conditions where the possible heterogeneous loss of the active species was negligible. All the relevant species were detected mass spectrometrically at their parent peaks: m/e ) 35 (Cl+), 51 (ClO+), 67 (OClO+), 70 (Cl2+), 127 (I+), 143 (IO+), 254 (I2+). The purities of the gases used were as follows: He > 99.9995% (Alphagaz) was passed through two liquid nitrogen traps; NO > 99.0%, was purified by trap-to-trap distillation to remove NO2 traces; NO2 > 99.00% (Alphagaz); O2 > 99.995% (Alphagaz); I2, ultrapure Normatom (Prolabo); CF3I > 99% (Fluorochem); OClO >95%; Cl2 > 99% (Ucar); Br2 > 99.99% (Aldrich). Results Reaction I + OClO (2). The measurements of the rate constant for the reaction
I + OClO f IO + ClO
(2)
were conducted in excess of iodine atoms over OClO in order to avoid complications due to the formation of I atoms as the products which may be formed in the fast secondary reaction
IO + IO f products (12) 3 -1 -1 7 k12 ) 1.7 × 10 exp(1020/T) cm molecule s -12
There is no recommendation so far for the branching ratios of reaction 12, but the channel forming 2I + O2 is very likely by comparison with the reactions ClO + ClO or BrO + BrO.7
15132 J. Phys. Chem., Vol. 100, No. 37, 1996
Bedjanian et al. influence of the OClO reformation on the experimental values of k2 was lower than 3%. Thus, the impact of reaction 1a could be considered as negligible. A second series of experiments were carried out at room temperature, with the detection of the products of reaction 2, IO and ClO. Quantitative measurements were made for ClO radicals but not for IO radicals, since complications could arise from the fast IO self-reaction for which the channels are not well established so far. Under the same experimental conditions, the kinetics of both OClO decay and ClO formation were monitored. It has to be noted that OClO contributed to the ClO mass spectrometric signal (at m/e ) 51) due to fragmentation in the ion source, although this latter operated at relatively low electron energy (25 eV). This contribution of OClO at m/e ) 51 was quantified by flowing a known concentration of OClO into the reactor, which allowed for correcting the ClO+ kinetic curves. As already mentioned, the ClO absolute concentrations were measured using the titration reaction 8 with an excess of Cl. This calibration method was preferred to that using the titration reaction
Figure 2. Reaction I + OClO (2): examples of pseudo-first-order plots at T ) 353 K (×) and 300 K, obtained from the monitoring of OClO consumption ([) and ClO production (]).
The configuration used for the introduction of the reactants into the reactor is shown in Figure 1a. I atoms, produced in reaction 4 in excess of I2 ([I2] ) (1-2) × 1014 molecule cm-3) were flowed into the reactor through the movable injector whereas OClO molecules were introduced through the sidearm tube of the reactor. The experimental conditions were as follows: total pressure ) around 1 Torr, temperature range ) 288-353 K, flow velocities ) 550-700 cm s-1; initial concentrations of OClO ) (5-10) × 1011 molecule cm-3, [I] ) (1.2-18.2) × 1013 molecule cm-3. In each run, the iodine atom concentration was constant along the reactor (the measured heterogeneous loss was less than 1 s-1 over the temperature range of the study). However, measurements were not reliable at T < 288 K, where formation of I2 crystals in reactor was observed. The kinetics of the OClO consumption in reaction 2 were monitored by mass spectrometry, and the reaction rate constant was derived from the expression
-d[OClO]/dt ) k2′[OClO] with k2′ ) k2[I] Examples of pseudo-first-order plots for OClO decay are given in Figure 2. The linear plots obtained have a zero intercept, which is expected from the observation of a negligible heterogeneous loss of OClO. All the results obtained as well as the corresponding experimental conditions are given in Table 1. Figure 3 shows the temperature dependence of k2. The resulting Arrhenius expression is
(13)
k13 ) 6.4 × 10-12 exp[-(290 ( 100)/T] cm3 molecule-1 s-1 7 since the uncertainty on k2 was reduced when the ratios [ClO]/ [OClO] were considered from the titration experiments using reaction 8, which did not require the absolute concentrations of the two species. The value of k2 was derived from modeling calculations of the ClO experimental kinetics, using the same reaction set as described above: reaction 1 (channel specification is not important), reaction 2 (assuming a unity branching ratio for the IO + ClO forming channel), reaction 12 (without channel specification), and wall losses for IO and ClO. Reaction 3 is too slow, as shown below, and can be disregarded. The k2 values thus obtained are reported in Table 1 and illustrated in Figure 2. These values are in excellent agreement with those obtained from the OClO consumption kinetics. It has to be noted that, in the above modeling, the fits for k2 were not sensitive to the other reaction rate constants: when all reactions were ignored (except reaction 2), the difference between the uncorrected and corrected value of k2 ranged from 0% to 20% for the concentration of iodine atoms ranging from 2.3 × 1013 to 1.1 × 1014 molecule cm-3, respectively. An example of kinetics of OClO consumption and ClO formation, fitted with the same value of k2 is also shown in Figure 4. In conclusion, these results show that reaction 2 proceeds exclusively, in the temperature range of the study through one channel forming IO and ClO:
I + OClO f IO + ClO
(2)
This precludes the occurrence of the other exothermic channel, which would have required a complex rearrangement:
k2 ) (9.7 ( 3.0) × 10-12 exp[-(1190 ( 200)/T] cm3 molecule-1 s-1 (288 K < T < 353 K) with twice standard deviation for the activation factor, E/R. To check the possible influence of the reverse reaction 1a:
ClO + IO f I + OClO
ClO + NO f Cl + NO2
(1a)
on the determination of k2, a simulation has been made with the simple chemical system including reactions 1a, 2, and 12 and the wall losses of ClO (3 s-1) and IO (50 s-1) (these wall loss rates are experimental values). Assuming a branching ratio of unity for channel 1a of reaction 1, it was observed that the
I + OClO f ICl + O2 ∆H ) -43.5 kcal mol-1 (2′) Reaction I + ClO (3). The two possible exchange channels of this reaction are
I + ClO f Cl + IO ∆H ) 5.7 ( 0.9 kcal mol-1 (3) f O + ICl ∆H ) 14.3 kcal mol-1
(3′)
A low rate constant is expected, in the temperature range of the study, for these two channels which are significantly
Kinetics of Reactions of I and of Cl Atoms
J. Phys. Chem., Vol. 100, No. 37, 1996 15133
TABLE 1: Reaction I + OClO (2): Experimental Conditions and Results 288 K [I]a 2.1 4.5 6.4 6.4 11.3 14.5 18.2
k2′, s-1 b 2.8 6.7 9.0 10.4 15.4 22.4 27.3
1.51 ( 0.12d
300 K
310 K
k2′, s-1
[I] 2.3 3.3 5.7 6.8 6.9 8.3 8.6 10.5 11.0 12.5
(3.8)c
4.5 7.2 (6.0) 9.2 (10.2) 12.6 13.7 17.2 15.3 (17.5) 16.8 (22.0) 19.7 24.6
1.90 ( 0.23
320 K
330 K
343 K
353 K
[I]
k2′, s-1
[I]
k2′, s-1
[I]
k2′, s-1
[I]
k2′, s-1
[I]
k2′, s-1
2.6 3.6 5.0 6.0 6.1 8.0 9.6 10.7 12.4
5.7 8.1 9.9 13.3 12.1 16.0 20.5 20.6 27.5
1.4 2.1 2.9 4.7 5.2 5.6 6.0 8.0 8.1 8.8 10.8 11.0 14.0 14.0
4.1 6.1 8.2 10.0 13.7 11.6 15.4 17.6 19.6 20.8 25.0 27.8 36.4 37.4
1.9 3.1 4.9 6.2 8.5 9.8 10.0 11.2 12.0 15.5
4.05 7.25 12.05 14.4 21.9 24.7 27.1 29 26.8 3.7
1.2 2.1 2.9 4.9 7.8 9.6
2.8 6.6 10.0 13.9 24.1 28.6
1.8 2.7 2.9 4.4 4.7 4.8 5.7 6.3 7.3 9.7 9.9 10.6
5.3 9.9 9.6 17.1 16.3 14.9 19.6 21.4 22.3 34.2 33.9 35.4
2.09 ( 0.23
2.55 ( 0.22
2.44 ( 0.25
3.0 ( 0.25
3.39 ( 0.26
Concentrations are in 1013 molecule cm-3. b k2′ ) k2[I]. c In parentheses: results obtained from modeling of ClO formation kinetics (see text). d Rate constants (k ) in 10-13 cm3 molecule-1 s-1 units; the error is twice the standard deviation. 2 a
Figure 3. Reaction I + OClO (2): temperature dependence of the rate constant k2 ) (9.7 ( 3.0) × 10-12 exp{-1190 ( 200)/T} cm3 molecule-1 s-1.
endothermic. The experiments were carried out by monitoring the ClO decay in excess of iodine atoms. Again, reaction 8 was used to generate ClO radicals. Special attention was borne to use of a chemical system free of both Cl and OClO. Otherwise, the fast reverse reaction 4 or the reaction between I and OClO (reaction 2) would have reproduced ClO, leading to errors in the k3 measurements. Therefore, the configuration for the introduction of the reactants, shown in Figure 1a, has been used in these experiments. OClO was introduced through the outer tube of the movable injector and was converted into ClO by an excess of Cl atoms flowed through the central tube. Excess Cl atoms were further rapidly consumed by reaction with excess I2 in the main reactor to give iodine atoms, the other reactant of reaction 3. Thus, only the I and ClO reactants were present in the reactor together with I2, Cl2, and ICl as unreactive species. Experiments were performed at P ≈ 0.9 Torr and at two temperatures: T ) 297 and 353 K. The maximum concentration used for iodine atoms was 2.25 × 1014 molecule cm-3 at T ) 297 K and 9.6 × 1013 molecule cm-3 at T ) 353 K, with [ClO]0 ) (7-8) × 1011 molecule cm-3 at these two temperatures. It was observed that an increase of I concentrations did not change the measured pseudo-first-order rate
Figure 4. Reaction I + OClO (2): examples of experimental (points) and simulated (solid lines) data for OClO consumption and ClO production at P ) 1 Torr, T ) 300 K, and [I] ) 3.3 × 1013 molecule cm-3 ([) and 8.6 × 1013 molecule cm-3 (b). The best fit was obtained for k2 ) 1.9 × 10-13 and 1.8 × 10-13 cm3 molecule-1 s-1, respectively.
constant for the ClO decay, within the experimental uncertainty range. The rate of ClO decay was the same as the ClO wall loss (measured in absence of I atoms): (3.0 ( 0.8) s-1 at T ) 297 K and (2.5 ( 0.8) s-1 at T ) 353 K. From these data, the following upper limits of k3 were obtained:
k3 e 7 × 10-15 cm3 molecule-1 s-1 at T ) 297 K k3 e 1.7 × 10-14 cm3 molecule-1 s-1 at T ) 353 K Reaction Cl + IO (4). In the study of this reaction, the experiments were carried out by monitoring the IO decay in the presence of an excess of Cl atoms. The configuration used for the introduction of the reactants into the reactor is shown in Figure 1b. Cl atoms were produced in the microwave discharge from highly diluted Cl2 in helium and were introduced through the internal tube of the movable injector. Oxygen atoms produced in the second microwave discharge were flowed through the outer tube of the injector and reacted with I2 or CF3I to yield IO radicals in the main reactor.
15134 J. Phys. Chem., Vol. 100, No. 37, 1996
Bedjanian et al.
The major experimental complication in this study came from the occurrence of side reactions of Cl atoms with the IO precursor, either I2 (reaction 5) or CF3I:
Cl + CF3I f CF3 + ICl -10
k14 ) 1.62 × 10
(14) -1 -1 13
3
exp(-2056/T) cm molecule
s
It has to be noted that the rates of the Cl consumption observed in the presence of either I2 or CF3I were significantly higher than one would expect from the existing rate constants for reactions 5 and 14, respectively. For reaction 14, there is no clear explanation for this difference. For reaction 5, we have remeasured the rate constant and the value obtained for k5 has been found to be about 1 order of magnitude higher than the only determination reported in the literature14 (see below). Two series of experiments, using the two sources of IO radicals, were carried out to study reaction 4. In the first one, reaction 9 was used to produce IO and the experimental conditions were as follows: temperature range ) 290-355 K, total pressure ) 0.9-1.0 Torr, linear flow velocities ) 12402130 cm s-1, concentration of molecular oxygen (precursor of O atoms) ) (1-3) × 1011 molecule cm-3, [I2] ) (0.7-2.0) × 1012 molecule cm-3. The initial concentrations of IO radicals were in the range (5-10) × 1010 molecule cm-3. In all the experiments, both IO radicals and Cl atoms were monitored. It was observed that the concentration of Cl atoms decreased very rapidly in the first part of the reaction zone. This was due to the fast reaction of Cl with I2 (reaction 5), and a simultaneous disappearance of I2 was also detected. Then, a much slower decay of Cl was observed, which was caused by the secondary reaction of Cl with ICl produced in reaction 5:
Cl + ICl f I + Cl2 -11
k15 ) 1.0 × 10
3
(15) -1 -1 15,16
cm molecule
s
The IO kinetics were measured in this second reaction zone and k4 could be derived from the expression
[IO] ) [IO]0 exp(-k4′t) with k4′ ) k4[Cl] + kw [Cl] is the mean value of Cl concentration along the reaction zone (see below), and kw is the heterogeneous loss rate of IO. The walls of the reactor were coated with halocarbon wax by two different methods. In the first one, the halocarbon wax was deposited by an evaporation-condensation technique. A degradation of this coating could be observed. and the wall loss rate measured for IO ranged from 30 to 100 s-1. This wall loss was found to decrease with increasing temperature. The second method used a direct melting of the wax on the surfaces of the reactor. In this case, the IO wall loss rate was less than 10 s-1 for the whole temperature range of the study. It has to be noted that the results obtained for the measurement of k4 did not depend on the “quality” of the reactor surface, and in all the cases, the zero intercepts of the pseudo-first-order plots (k4′ as a function of [Cl]) were found in good agreement with the values of kw(IO) measured directly. The experimental conditions and results are summarized in Table 2. The change in Cl concentration along the reaction zone used for the kinetic measurements was in the range 1030%. In separate experiments, it was shown that the increase of I2 initial concentration from 1.0 × 1012 up to 6 × 1012 molecule cm-3 did not affect the measured values for k4 although the Cl concentration was changed significantly. To check if the use of the mean concentration of Cl atoms did not affect the calculation of k4 given above, k4 was also derived by fitting the experimental IO kinetics from a simulation
TABLE 2: Reaction Cl + IO (4): Experimental Conditions and Results N/expt
T (K)
[Cl] (1012 molecule cm-3)
k4 (10-11 cm3 molecule-1s-1)a
8 9 8 7 8 8 10 12
290 296 300 310 320 325 335 355
0.3-6.1 0.3-2.8 0.5-7.2 0.6-5.4 0.8-7.7 1.5-7.9 0.5-8.2 1.2-11.3
4.9 ( 1.0 4.4 ( 0.7 4.1 ( 0.7 3.8 ( 1.0 4.5 ( 0.7 4.8 ( 0.9 4.5 ( 0.6 3.9 ( 0.8
a The error is one standard deviation with addition of a 10% uncertainty for [Cl] measurements.
using the observed Cl temporal profiles. The k4 values obtained by the two methods were in agreement within 5%. The pseudo-first-order rate constants of the IO decay were corrected for the axial diffusion of IO. The diffusion coefficient of IO in He, DIO -He) 0.53 × (T/296)1.66 atm cm2 s-1, was calculated from DXe -He.17 The correction was less than 4%, except for a few experiments at 335 and 355 K, where the correction could reach 8%. As shown in Table 2, the measured values for k4 are not temperature dependent, within the experimental uncertainty, in the temperature range of the study. Finally, the value of k4, in the range 290-355 K, is
k4 ) (4.4 ( 1.0) × 10-11 cm3 molecule-1 s-1 In a second series of experiments, reaction10 was used as the source of IO radicals and the experimental conditions were similar to those described above. The concentration of CF3I was in the range (1-3) × 1013 molecule cm-3. The values of the rate constant k4, measured using the same method as in the first series of experiments, were found to be lower by approximately 50% over all the temperature range. This difference was explained by an incomplete consumption of O atoms at the CF3I concentrations used within the residence time (0.01s) in the IO production zone. Thus, in the reaction zone, IO could be produced at a rate comparable to that of the IO consumption in reaction 4, leading to an underestimation of k4. This complication could not be suppressed since the use of higher concentrations of CF3I would have led to a significant consumption of Cl atoms by reaction 14. It has to be mentioned that such a complication did not exist in the first series of experiments where a very fast consumption of I2 by Cl took place. In this case, the re-formation of IO was also possible via the reaction of O with ICl produced by the reaction of I2 with Cl:
O + ICl f Cl + IO ∆H ) -(8.6 ( 0.9) kcal mol-1 (16) The rate constant of reaction 16 is not known, but this reaction was shown to play no role here, since the variation of I2 concentration did not affect the measured value of k4. Finally, the kinetics of IO consumption in the second series of experiments were simulated taking into account the occurrence of reaction 10 in the reaction zone. The initial concentrations of O atoms used in these calculations were about (2-3) × 1010 molecule cm-3. The values of k4 derived from the fitting procedure of experimental IO kinetics, at different temperatures ranging from 285 to 365 K, were in the range (3.5-5.0) × 10-11 molecule cm3 s-1. These values are in fair agreement with those obtained in the first series of experiments. Considering the uncertainty on k4, which mainly resulted from the uncertainty on O atom concentration, the preferred value of k4 is that
Kinetics of Reactions of I and of Cl Atoms
J. Phys. Chem., Vol. 100, No. 37, 1996 15135 was negligible in most experiments. However, in a few cases, the Cl loss could reach 30% over the reaction distance. Experimental data are summarized in Table 3 and Figure 5. k5 was found to be temperature independent in the range of the study (296-365 K). The linear fit to all the data yields the value for k5:
k5 ) (2.1 ( 0.3) × 10-10 cm3 molecule-1 s-1 The error is one standard deviation plus 10% due to the uncertainty on the concentrations of Cl atoms. Discussion
Figure 5. Reaction Cl + I2 (5): pseudo-first-order plots at T) 296, 315, 340, and 365 K.
TABLE 3: Reaction Cl + I2 (5): Experimental Conditions and Results N/expt
T (K)
[Cl] (1012 molecule cm-3)
k5 (10-10 cm3 molecule-1 s-1)
9 10 8 10
296 315 340 365
0.4-4.0 0.46-3.9 0.4-3.2 0.47-5.3
2.07 ( 0.05 2.20 ( 0.08 2.07 ( 0.07 2.02 ( 0.07
obtained in the first series of experiments: k4 ) (4.4 ( 1.0) × 10-11 cm3 molecule-1 s-1 (T ) 290-355 K). Reaction 4 has two possible channels:
Cl + IO f I + ClO ∆H ) -(5.7 ( 0.9) kcal mol-1 f O + ICl ∆H ) 8.6 ( 0.9 kcal mol-1
(4) (4′)
However, channel 4′ is too endothermic to proceed with a significant rate in the temperature range of the study. Thus, the value measured for k4 very likely corresponds to the channel forming I atoms and ClO radicals. It has to be noted that the secondary reaction of IO with ClO as well as the IO selfcombination was negligible and thus did not affect the measurements of k4 at the low initial concentrations used for IO. Reaction Cl + I2 f I + ICl. The production and introduction of the reactants into the reactor are shown in Figure 1b. Cl atoms, produced in the microwave discharge of a Cl2/ He mixture, were introduced into the reactor through the central tube of the movable injector. Molecular iodine was flowed through the side-arm tube. Experiments were conducted under pseudo-first-order conditions with an excess of Cl atoms over I2 molecules. The experimental conditions were as follows: temperature range ) 296-365 K, total pressure about 1 Torr, flow velocity ) 2000-2500 cm s-1, initial I2 concentration ) (1-2) × 1011 molecule cm-3. The ranges of Cl concentrations used and the k5 values obtained from I2 kinetics are listed in Table 3. The pseudo-first-order plots are also given in Figure 5: -d(ln[I2])/dt ) k5′ ) k5[Cl]. The experimental values of k5′ were corrected for the axial diffusion of I2. This correction was lower than 6%, using diffusion coefficients for I2 in He derived from DXe-He:17 DI2-He ) 0.4, 0.44, 0.5, and 0.56 atm cm2 s-1 at T ) 296, 315, 340, and 365 K, respectively. The wall loss of Cl, kw(Cl) ) 8 ( 1 s-1, as well as the Cl consumption due to the reactions of Cl with I2 and with ICl,
The kinetics of the reactions 2-4 have been studied for the first time. The only previous study of reaction 514 provided an indirect determination of k5 ) (2.4 ( 1.0) × 10-11 cm3 molecule-1 s-1 at T ) 298 K. The present value of k5, obtained by a more direct way, is about 1 order of magnitude higher than that of this earlier work.14 A higher value for k5 is acceptable when compared with the rate constants measured for similar reactions of molecular iodine with other atoms: H + I2, k) 6.6 × 10-10 exp(-20/T) cm3 molecule-1 s-1;10 O + I2, k ) 1.4 × 10-10 exp((0 ( 250)/T) cm3 molecule-1 s-1;7 F + I2, k ) (4.3 ( 1.1) × 10-10 cm3 molecule-1 s-1(T ) 298 K).18 The kinetic and mechanistic data obtained for the reactions 2-4 in the present work give some new information about the enthalpy of formation of IO, ∆Hf(IO), which is presently not well established. First, an upper limit of ∆Hf(IO) can be obtained considering that the activation energy measured for k2 is an upper limit for the enthalpy of reaction 2. Therefore, ∆Hf(IO) e (2.38 ( 0.4) - ∆Hf(ClO) + ∆Hf(I) + ∆Hf(OClO) (units are kcal mol-1). Using the recommended values of ∆Hf7 (∆Hf(ClO) ) 24.4 kcal mol-1, ∆Hf(I) ) 25.52 kcal mol-1, ∆Hf(OClO) ) 22.6 ( 1 kcal mol-1), the calculated value, ∆Hf(IO) 26.1 ( 1.4 kcal mol-1, leads to the upper limit
∆Hf(IO) e 27.5 kcal mol-1 Second, the kinetic data obtained in the present study for the forward (4) and reverse (3) reactions
Cl + IO T I + ClO
(4,3)
allow for the determination of a lower limit for ∆Hf(IO). The enthalpy of reaction 4 can be calculated from the equilibrium constant
K ) k4/k3 ) exp{-(∆Hr - T∆Sr)/RT} Using the values obtained in the present work at 298 K, k3 e 7 × 10-15 and k4 ) 3.4 × 10-11 cm3 molecule-1 s-1, which is the lower limit of k4 ) (4.4 ( 1.0) × 10-11 cm3 molecule-1 s-1, the value derived for K is k4/k3 > 4860. This value can be used in combination with thermochemical data to calculate ∆Hr and subsequently, ∆Hf(IO). Using the heats of formation of species from ref 7, the entropy data S°298(Cl) ) 39.48 cal K-1 mol-1, S°298(I) ) 43.21 cal K-1 mol-1, and S°298(ClO) ) 54.17 cal K-1 mol-1 from ref 19, and S°298(IO) ) 57.12 cal K-1 mol-1,20 it results that ∆Hr e -4.8 kcal mol-1 and ∆Hf(IO) g 25.8 kcal mol-1. By combining this result with the upper limit of ∆Hf(IO) calculated above, we obtain
25.8 kcal mol-1 e ∆Hf(IO) e 27.5 kcal mol-1
15136 J. Phys. Chem., Vol. 100, No. 37, 1996
Bedjanian et al.
∆Hf(IO) may be also estimated from the kinetic information presently available for the forward and reverse reactions of the equilibrium
I + OClO T IO + ClO
(2,1a)
The enthalpy of reaction 2 can be calculated from the expression for the equilibrium constant as above. However, only a lower limit of this equilibrium constant can be presently calculated since only the overall rate constant of the reaction between IO and ClO has been measured so far.8 Using the values at 298 K, k2 ) 1.9 × 10-13 cm3 molecule-1 s-1 and k1 ) 1.8 × 10-11 cm3 molecule-1 s-1,8 and the known thermochemical parameters,7,19,20 it can be calculated that ∆Hr e 4.7 kcal mol-1 and ∆Hf(IO) e 29.4 kcal mol-1. By varying the branching ratio for channel 1a of reaction 1 from 1% to 100%, the values of ∆Hr range from 1.9 to 4.7 kcal mol-1, and those of ∆Hf(IO), from 25.6 to 28.4 kcal mol-1. An additional error of 1 kcal mol-1 arises from the uncertainty on ∆Hf(OClO). This gives 24.6 kcal mol-1 e ∆Hf(IO) e 29.4 kcal mol-1. This range, which is not inconsistent with the range 25.8-27.5 kcal mol-1 given above, will be better defined when the branching ratios for reaction 1 become available. By comparison with the ClO + BrO reaction, where the similar channel (ClO + BrO f Br + OClO) has a branching ratio of 0.5,7 the equilibrium constant is 2.1 × 10-2, giving ∆Hr ) 4.25 kcal mol-1 and ∆Hf(IO) ) 28 ( 1 kcal mol-1. Finally, the preferred value of ∆Hf(IO) from the present work is 25.8 kcal mol-1 e ∆Hf(IO) e 27.5 kcal mol-1, which can be expressed as
∆Hf(IO) ) 26.7 ( 0.9 kcal mol-1 This result appears to be consistent with the upper limit of Gilles et al.21 (∆Hf(IO) e 29.5 kcal mol-1) obtained from the study of the IO + CH3 forming channel of the reaction between O(3P) and CH3I. The present result also agrees with the most recent recommendations from the NASA evaluation panel (∆Hf(IO) e 28 kcal mol-17) and from the IUPAC panel (∆Hf(IO) ) 25.6 kcal mol-1 22). This latter is from Reddy et al.23 However, our result is significantly lower than the value extracted from ref 24 (∆Hf(IO) ) 31 kcal mol-1). Acknowledgment. This work has been performed within the Environment Program of the European Commission (LEXIS project).
References and Notes (1) Zafiriou, O. C. J. Geophys. Res. 1974, 79, 2730. (2) Chameides, W. L.; Davis, D. D. J. Geophys. Res. 1980, 85, 7383. (3) Jenkin, M. E.; Cox, R. A.; Candeland, D. E. J. Atmos. Chem. 1985, 2, 359. (4) Chatfield, R. B.; Crutzen, P. J. J. Geophys. Res. 1990, 95, 22319. (5) Barnes, I.; Bonsang, B.; Brauers, T.; Carlier, P.; Cox, R. A.; Dorn, H. P.; Jenkin, M. E.; Le Bras, G.; Platt, U. Air Pollution Research Report. OCENO-NOX Project; Commission of the European Communities, Brussels, 1991. (6) Solomon, S.; Garcia, R. P.; Ravishankara, A. R. J. Geophys. Res. 1994, 99, 20491. (7) De More, W. D.; Sander, S. P.; Golden, D. M.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R.; Kolb, C. E.; Molina, M. J. Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling; NASA, JPL, California Institute of Technology: Pasadena, CA, 1994. (8) Ravishankara, A. R. Presented at the International Conference on Ozone in the Lower Stratosphere, Halkidiki, Greece, 1995. (9) Lancar, I. T.; Laverdet, G.; Le Bras, G.; Poulet, G. J. Phys. Chem. 1990, 94, 278. (10) Lorenz, K.; Wagner, H. Gg.; Zellner, R. Ber Bunsen-Ges. Phys. Chem. 1979, 83, 556. (11) Nicovich, J. M.; Wine, P. H. Int. J. Chem. Kinet. 1990, 22, 379. (12) Herron, J. T. J. Phys. Chem. Ref. Data 1988, 17, 967. (13) Ahonkhai, S. I.; Whittle, E. Int. J. Chem. Kinet. 1984, 16, 543. (14) Kuznetsova, S. V.; Maslov, A. I. Khim. Fiz. 1987, 6, 1554. (15) Clyne, M. A. A.; Cruse, H. W. J. Chem. Soc., Faraday Trans. 2 1972, 68, 1377. (16) Chesnokov, E. N. Khim. Fiz. 1991, 10, 204. (17) Marrero T. R.; Mason, E. A. J. Phys. Chem. Ref. Data 1972, 1, 3. (18) Appelman, E. H.; Clyne, M. A. A. J. Chem. Soc., Faraday Trans. 1 1975, 71, 2072. (19) JANAF. JANAF Thermochemical Tables, 3rd ed.; National Bureau of Standards: Washington, D.C., 1985. (20) Rayez, M. T. Private communication. (21) Gilles, M. K.; Turnipseed, A.; Burkholder, J. B.; Ravishankara, A. R. Presented at the International Conference on Ozone in the Lower Stratosphere, Halkidiki, Greece, 1995. (22) Atkinson, R.; Baulch, D. L.; Cox, R. A.; Hampson, R. F.; Kerr, J. A.; Troe, J. J. Phys. Chem. Ref. Data 1992, 21, 1125. (23) Reddy R. R.; Rao T. U. R.; Reddy A. S. R. Indian J. Pure Appl. Phys. 1989, 27, 243. (24) Laszlo, B.; Kurylo, M. J.; Huie, R. E J. Phys. Chem. 1995, 99, 11701.
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