J . Phys. Chem. 1990, 94, 3860-3863
3860
Physicochemical Properties of Chlorine Oxides. 1. Composition, Ultraviolet Spectrum, and Kinetics of the Thermolysis of Gaseous Dichlorine Hexoxide Maria I. L6pez and Juan E. Sicre* Instituto de Inuestigaciones Fisicoquimicas TeBricas y Aplicadas (INIFTA). Facultad de Ciencias Exactas, Unicersidad Nacional de La Plata, Casilla de Correo 16, Sucursal.4, (1900) La Plata, Argentina (Receiced: April 17, 1989; In Final Form: November 7 , 1989)
The experimental evidence presented here shows that gaseous dichlorine hexoxide at room temperatures has essentially the composition C1206and not CIO, as currently accepted. Vapor pressure measurements of liquid C1206between 5 and 45 O C have been made. The absorption cross section values of gaseous C1206are approximately 3 times higher than those assigned to CIO, by other authors. The products of the thermal decomposition of gaseous C1206are mainly chlorine perchlorate and oxygen. The kinetics of the decomposition over the C1206pressure range 0.5-1.2 Torr, at 20-30 OC,follows a first-order rate equation and is independent of the total pressure. A unimolecular process, namely, C1206 Cl2O, + 02,proceeding exp[(-10370 at the high-pressure limit through a complex fission reaction is postulated. The rate constant k , = 10(".4*'.2) 8 0 0 ) / q s-I is determined.
-
Introduction Although dichlorine hexoxide is a compound known for about one and a half centuries,' its molecular composition, structure, and other properties are still unclear. Undoubtedly these uncertainties arise primarily from the instability of this substance at room temperature. In 1925 Bodenstein et al.z established the ratio of oxygen to chlorine atoms in dichlorine hexoxide as three to one and they also determined the molecular weight of the compound in solution in agreement with the formula C1206. Later Goodeve and Todd3 reported that this oxide exists in the gaseous state entirely as CIO,. Goodeve et aL4 found a trioxidehexoxide equilibrium in the liquid phase with a very low dissociation energy (1.7 kcal/mol). Later on, they obtained the visible-ultraviolet spectra of the liquid and the gasSand determined the vapor pressure of the solid and liquid.6 Recently, Schack and Christe7 performed an infrared matrixisolation study which discounts the formation of CIO, through CI2O6dissociation in the condensed phase. More recently, an oxygen-bridged structure 02C10C103 for the gaseous oxide at room temperature has been proposed by Jansen et al.839based on an infrared spectroscopic study. They did not find any evidence for the presence of monomeric ClO,. Dichlorine hexoxide and, in good yield, chlorine perchlorate have been produced in the C102 photolysis at 436 nm.Io No exact chlorine balance was obtained, however, by applying the available spectrometric data for the system Clz06e=$ 2C1035and for CIOC103.11 The present study was undertaken in order to determine some properties of gaseous CI2O6,in particular its molecular composition, ultraviolet spectrum, and thermal stability.
Experimental Section A conventional apparatus for static gas kinetics studies was utilized. Reservoirs, tubes, and stopcocks that were in direct contact with C1206 were made of fused quartz. A 100-cm3 spherical vessel was used as the reaction vessel. Pressure measurements were made with a MKS (Baratron) pressure transducer connected to a chart recorder. Minimal amounts of Halocarbon ( 1 ) Millon, E. Ann. 1843, 46, 281 (p 312).
(2) Bodenstein, M.; Harteck, P.; Padelt, E. Z . Anorg. Chem. 1925, 147, 233. (3) Goodeve, C. F.; Todd, F. A. Nature 1933, 132, 514. (4) Farquharson, J.; Goodeve, C. F.; Richardson, F. D. Trans. Faraday SOC.1936, 32, 790. ( 5 ) Goodeve, C. F.; Richardson, F. D. T r a m . Faraday Soc. 1937,33,453. ( 6 ) Goodeve, C. F.; Richardson, F. D. J . Chem. SOC.1937, 294. (7) Schack, C. J.: Christe, K . 0. Inorg. Chem. 1974, 13, 2378. (8) Jansen, M.; Tobias, K . M.; Willner, H. Natunvissemchaften 1986, 73, 734. (9) Jansen, M.; Schatte, G.; Tobias, K. M.; Willner, H . Inorg. Chem. 1988. 27. 1703. ( I O ) Lbpez, M . 1.; Sicre, J. E.; Schumacher, H. J., to be published. ( 1 1 ) Lbpez. M. 1.; Sicre, J. E. J . Phys. Chem. 1988, 92, 563 and references therein.
0022-3654/90/2094-3860.$02.50/0
TABLE I 10
Pc12061PocIo* PC1206IPC12 PO2lPCI2
oc
20
0.48
O C
0.49 1.04 2.8
1 .o
2.8
30 "C 0.49
1.05 2.8
grease were used as lubricant for stopcocks and joints because C1206has a great tendency to dissolve in it. UV spectra were obtained with a Cary 14 spectrophotometer. The C1206gas samples were prepared in a IO cm long (2 cm diameter) quartz cell and fractional evaporations were made in a U-tube inserted between the cell and the stopcock. The temperature was kept constant (fO.l "C) by circulation of water from a thermostatic bath through an outer jacket on the UV cell. IR spectra were obtained with a Perkin Elmer 325 spectrophotometer using a Pyrex glass gas cell with AgCl windows. Chlorine dioxide was prepared by the reaction of potassium chlorate, oxalic acid, and dilute sulfuric acid;I2 it was dried with concentrated sulfuric acid and phosphorus pentoxide, purified by low-temperature distillation, and stored in the dark at liquid air temperature. Ozone was obtained from electrolysis of oxygen in a 15-kV ac electric discharge. The 6% ozone in oxygen mixture was purified by passage through a -100 "C cold trap before storage in a light-protected 2-L glass bulb. C1206was synthesized in the reaction vessel in order to minimize its own decomposition. The method of Schumacher and Stieger7*I3 turned out to be very useful for this purpose. C1206, free of impurities, was produced when measured amounts of C102 reacted quantitatively at 0 OC with excess ozone diluted with oxygen. Unreacted ozone and O2were eliminated by fractional evaporation at -50 OC. The reaction mechanism is CIO, + 0, C103 + O2 (1)
-
CIO,
+ C103
M
C1206
(2)
Nitrogen and oxygen were obtained from high-quality commercial cylinders and passed through a -183 "C cold trap before use. Results and Discussion Gas-Phase Composition, Vapor Pressure, and Decomposition Products. A series of experiments were performed at IO, 20, and 30 "C. For each run, C1206,prepared in the reaction vessel, was allowed to warm up at the selected temperature. In all cases, fast pressure increases folowed by much slower yet practically uniform ( 1 2) Brauer, G. Handbook of Preparative Inorganic Chemistry; Academic Press: New York, 1963. ( 1 3 ) Schumacher, H. J.; Stieger, G . Z. Anorg. Chem. 1929, 184, 272.
0 1990 American Chemical Society
Physicochemical Properties of Chlorine Oxides
/
The Journal of Physical Chemistry, Vol. 94, No. 9, 1990 3861
030'C 0
L
c 0
n a
0
Figure 1. Initial pressure of chlorine dioxide p0cIo2vs resultant pressure of dichlorine hexoxide pc1206, 0 , 0 , and 0, experiments at 10, 20, and 30 "C, respectively.
ones were observed. The first increase corresponds to evaporation while the second corresponds to the decomposition of the oxide. The initial oxide pressure was obtained by extrapolation to zero time from a pressure-time plot. The amount of dichlorine hexoxide formed, pc1206, vs chlorine dioxide pressure consumed, pocIo2,is plotted in Figure 1 at the three temperatures. In the cases where p,-1206is smaller than the equilibrium vapour pressure at a given temperature, the plots are straight lines passing through the origin, with slopes of 0.5. Upon reaching the equilibrium vapour pressure, pCIIo6 becomes independent of the CI02 initial pressure. For p , - l t ~below d the equilibrium vapor pressure, an analysis of the gaseous sample was made by complete thermal decomposition of the formed oxide. The resultant oxygen, PO*,was eliminated by evacuation at -196 OC and the residual chlorine pressure pCl2determined. The obtained pressure ratios are shown in Table I. From these values it can be concluded that within the experimental error (less than i 4 % ) the oxide in the gas phase has the composition C1206. From the pressuretime curves it was possible to prove that the rate of decomposition of the oxide increases as the temperature rises, although above 40 OC the zero time extrapolations were very inaccurate. Nevertheless, this rate is independent of the amount of liquid present in agreement with the results of Schumacher et a1.13 A log p vs 1/ T plot for C1206,between 0 and 45 OC, is presented in Figure 2. For comparison purposes the results given by Goodeve et aL6 are also depicted. The reported values at 10,20, and 30 OC (eight measurements at each temperature) can be represented by the expression (2729 f 90) log p = + (9.34 f 0.31) T from which a Trouton constant of 29.5 eu is calculated. This rather high value appears reasonable since CI2O6is highly selfassociated in the liquid phase. Moreover, the molecule should be strongly polarized across the oxygen-bridged structure 02C1+OCI03-7and it is readily transformed into an ionic solid.15 Experiments were done at temperatures between 20 and 57 OC in order to determine the products of the C1206 spontaneous decomposition. The reactant initial pressure was selected according to the temperature and in some cases up to 200 Torr of nitrogen or oxygen was added. The analysis of the products resulting after more than 95% decomposition of the reactant, which takes place in about 8 h at 20 OC and 8 min at 57 "C, was performed by IR (14) Arvia, A. J.; Basualdo, W. H.; Schumacher, H. J. Z.Anorg. Chem. 1956, 286, 58. (15) Pavia, A. C.:Pascal, J. L.; Potier, A. C.R. Acad. Sc. 1971, 0 7 2 , 1495. Tobias, K . M.; Jansen, M. Angew. Chem. 1986, 98,994; 2.Anorg. Chem. 1987, 550, 16.
I 1
- 0.5
\
1 32
34
+
\
36
IOL I T
Figure 2. Vapor pressure of CI,O,: (0)liquid between 5 and 45 "C; ( 0 ) solid at 0 "C; (+) ref 6 values. For the calculation of the log p vs 1/T equation the values at 10, 20, and 30 "C were only considered.
and UV spectrometric techniques combined with fractional evaporation. The reaction products whose relative quantities depend on the experimental conditions are, by order of importance, chlorine perchlorate, oxygen, chlorine dioxide, chlorine, and minimum amounts of dichlorine heptoxide (less than 3%). In cases of C1206 decomposition without addition of an inert gas, the reaction product yields of C1204and C102 were 60% and 15%, respectively. C102 is formed mainly in the first stage of the reaction. In cases where about 10 Torr of C102 was initially added to the reactant no yield increase of C102 was observed, but the yield of C1204rose to 90%. Even this could be improved when together with C102,200 Torr of an inert gas was added. In this case, if no addition of C102 was made, no more than 75% of CI2O4 could be obtained. These facts indicate that C102 is mainly formed by CI2O4decomposition,probably on the wall of the reaction vessel, and CIOz inhibits the wall reaction. The C1204instabilty would explain why other a ~ t h o r s ~ have * ~ , ~notJ ~found this species as a reaction product; however, C1204 was found in the liquid-* and gas-phaseg C1206thermal decomposition. Therefore, C1206decomposes in the gas phase mainly according to the reaction stoichiometry CI2O6 = C1204 + 02,and there is also a small contribution of other reaction which will be discussed later. The Ultraviolet Spectrum of Gaseous C1206. The C1206relative absorption cross section was determined by preparing this substance directly in the quartz UV cell as indicated in the previous section. In order to determine absolute absorption cross sections, gaseous C1206/C102samples were prepared at 20 OC by mixing O3and a known excess of C102. The spectrum was obtained after a few minutes such that ozone was consumed to a pressure less than 0.01 Torr.I6 C102 absorption was subtracted by using calibrated C102 spectra. For these experiments, C1206pressures were varied between 0.3 and 0.8 Torr. The experimental error can be estimated as i5% (maximum-minimum values) but rises to i 1 0 % at wavelengths where C102 absorption must be subtracted. In Figure 3 the gaseous C1206 absorption cross section u vs wavelength is presented. For comparison purposes the spectrum attributed to gaseous C103 by Goodeve et aLs is also given. Their smaller u values might be attributable to C1206 decomposition (16) Wongdontri-Stuper, W.; Jayanty, R. K. M.; Simonaitis, R.; and Heicklein, J. J . Photochem. 1979, 10, 163.
Ldpez and Sicre
3862 The Journal of Physical Chemistry, Vol. 94, No. 9, 1990
e e
'. ,
D 4
200
m e
o
---__ LOO
300 WAVELENGTH
o
(nml
Figure 3. Absorption cross section of gaseous CI2O6. Dashed line: values obtained after subtracting C102absorption. Points: values attributed to gaseous CIOl (ref 5 ) .
100
200
300
timin)
(absorbanceof C1206)vs time (min) at 20 O C (right and lower scales), and 30 "C (left and upper scales). The initial pressure of C1206was between 0.5 and 1.0 Torr at 20 "C and between 0.6 and 0.9 at 30 O C . Experiments without adding gas: 0, 0. Experiments with oxygen added: (e) 3 Torr, ( 0 ) 10 Torr, (e, 0 , O), 200 Torr, (0) 250 Torr. With nitrogen added: ( 0 ,0 , Q, 0 ) 200 Torr. ( 0 )200 Torr of N, plus 10 Torr of (210, (this experiment was discarded for evaluation Figure 5. In A
purposes).
7
11
(Jansen et aL8s9obtained at rcom temperature a considerably lower value t l / 2 = 8 min.) The calculated Arrhenius activation energy is E, = 20600 cal and the rate constant is given by the expression
1
/
I
///
01
k =
05
03
07
Pc,,o, Figure 4.
Absorbance vs CI2O6pressure at 200, 240, and 280 nm.
and the higher relative absorption values at the shortest wavelengths could be due to C1204. The absorbance vs CI2O6 pressure at three wavelengths is presented in Figure 4. Beer's law is obeyed within experimental error. I n addition, no noticeable absorbance variation during a 20 O C temperature jump in the UV cell was observed. These observations support the conclusion that the samples consist of pure C1206. Jansen et aL9 have reported an absorption band with only one maximum at 268 nm. This discrepancy can be explained by assuming that a band of an impurity (probably residual ozone) hides the 240-nm minimum. Kinetics of the Gaseous C1206 Thermal Decomposition. This study was carried out at 20 and 30 O C . The reaction was followed by measuring C1206absorbance at 280 nm up to approximately 90% reactant consumption. The initial C1206pressure was varied between 0.5 and 1.2 Torr depending on the temperature, total pressure was changed by adding 3-250 Torr of oxygen or nitrogen (200 Torr), and in some experiments, IO Torr of C102was added. Figure 5 shows a In A (absorbance) vs time plot at both temperatures. In all cases linear relationships are obtained. However, in the experiments performed without inert gas addition, some irregularities are observed in the first part of the reaction. The first-order rate constants obtained from the slopes at 20 and 30 O C are kzooc = (6.5 i 0.6) X min-I and k300C= (20.9 f 1 . 3 ) X min-' independent of total pressure in the range 0.5-250 Torr. Therefore, the half-lives are tl12(20"C) = 107.0 min and t1,2(30"C) = 33.2 min
~0(11.4*1.2)
exp( - 1 0 3 7 7 800) s-I
(1)
A similar value has been obtained for the activation energy of C1206decomposition in the liquid phases6 The obtainment of a first-order rate constant suggests that the rate-determining step of the reaction is a unimolecular process. In particular, the unimolecular reaction C1206 C1204 + O2 (4) in its high-pressure limit would explain the experimental results. In this case, k4 would correspond to k,, the limiting highpressure rate constant. The CI2O6transition pressure," p l 1 2 at , which the ratio k,,/k, = ( k w I is the experimental first-order rate constant) can be estimated by the treatment of B e n s ~ n . ~ ~ ~ ~ ~ If one assumes that the C1206structure is similar to that of C1207, this procedure leads to p l l z = 0.01 Torr, in agreement with the experimental results. (For C1207,20k , = 4.5 X I O l 5 exp(16 560/'j3 s-l and p l 1 2E 1 Torr; for this characteristic pressure the calculated number of effective internal degrees of freedom (serf)at 298 K is serf = 2 / 3 statal; the same approxima!ion was employed to calculate p l 1 2for C1206.) The experimental A value is rather low and it could be supposed that reaction 4 is a complex fission reaction in which more than one bond is broken through the formation of a rigid cyclic transition statel7.l9
-
0
cl' + \ / 'o 0
C/O
02
(4')
(17) Robinson, P. J.; Holbrook, K. A. Unimolecular Reactions, Wiley: London, 1972. ( 1 8 ) Benson, S. W. The Foundations of Chemical Kinetics; McGraw-Hill: New York, 1960. (19) Benson, S. W. Thermochemical Kinetics, 2nd ed.;Wiley: New York, 1976.
(20) Figini, R. V.; Coloccia, E.; Schumacher, H. J. 2. Phys. Chem. 1958, 14, 32.
The Journal of Physical Chemistry, Vol. 94, No. 9, 1990 3863
Physicochemical Properties of Chlorine Oxides followed by a fast isomerization of this chlorine oxide to the chlorine perchlorate stable structure. Other possible mechanisms have been discarded. Firstly, the dissociation into two radicals
-
C1206 followed by
M
C103 + C103
2CI03
(-2)
C1204 + O2
(3)
also leads to the experimental rate equation if reaction 3 is fast. However, assuming that the reaction mechanism for the C1206 formation from C102 and O3(reactions 1 and 2) is correct, reaction 3 cannot be fast. Secondly, a first-order rate equation with a pressure-independent rate constant can be obtained by combining reactions -2, 2, and 3 with the assumption that, even at the lower pressures studied, k,(M) >> k 3 . Therefore -d(Cl206)/dZ = +d(Cl,O,)/dt
= k3k_2/k2(C1206) =
k(C1206) However, this mechanism must be rejected because there is no experimental evidence for a clO3/c1206 equilibrium. A heterogeneous contribution for the C1206decomposition is possible in view of the large difference between the half-life of the reaction obtained from our experiments and that quoted by Jansen et al.899 The low A value experimentally obtained would support this reaction path. In terms of a Langmuir adsorption isotherm this rate constant would correspond to a catalytic reaction at low coverage of sorption sites.I9 Nevertheless, in our case the catalytic activity should be of minor importance since the firstorder rate constant is adhered to at pressures near the saturation point to more than 90% of reactant consumption. On the other hand, a strong inhibiting effect of chlorine dioxide on the catalytic decomposition of C1206 has been r e p ~ r t e d . ~ However, the rate observed in our experiments is unaffected by the presence of different amounts of C102.
With respect to the formation of C1207as a normal product of C1206decomposition, it could be explained through the following mechanism: a simple unimolecular fission reaction, that is, elimination of C102 by breaking the oxygen-bridged structure O,CI-O-CIO3
25 (2102 + clod
-
followed by the fast reactions 2C104 2C103 + O2 C104 + C103
M
(5)
(6)
C1207
(7) Reaction 5 can also account for the small C102 increase during the gas-phase C1206decomposition. The C1207yield increases with the temperature rise and this suggests a higher activation energy for reaction 5 compared to that for reaction 4 which was experimentally obtained. However, owing to the small contribution of reaction 5 and the instability of the chlorinated main reaction product (C1204),any inference about the kinetics of CI2O7 formation is uncertain. The reported data support the conclusion that “chlorine trioxide” in the gas phase at room temperature exists nearly completely associated in a dimeric form C1206. Therefore, the possible participation of C103 as a potential source for odd oxygen generation in the stratosphere2’ should be reformulated because previous data for this radical, especially that concerning its absorption cross section,22 are now known not to be applicable.
Acknowledgment. This research project was financially supported by the Consejo Nacional de Investigaciones Cientificas y Tecnicas and the Comisidn de Investigaciones Cientificas de la Provincia de Buenos Aires. We are indebted to Dr. C. C o b s and Dr. A. Croce for helpful discussions. Registry No. CI2O6,12442-63-6. (21) Handwerk, V.; Zellner, R . Ber. Bunsen-Ges. Phys. Chem. 1986, 90, 92. (22) ‘Chemical Kinetics and Photochemical Data for Use in Stratospheric Modeling”. NASA-JPL. Evaluation Number 7, 1985.
