Initiation Rate of LiC104 Decomposition
(4) (5) (6) (7) (8) (9) (10) (11)
(12) (13)
773
tion, by inversion of a temperature-dependent Laplace integral was demonstrated by S. H. Bauer, J. Chem. Phys., 6, 403 (1938); 7, 1097 (1939). S. W. Benson and H. E. O'Neai, Nat. Stand. Ref. Data Ser., Nat. Bur. Stand., No.21 (1970). H. Hartman, H. G. Bosch, and H. Heydtmann, Z. Phys. Chem. (Frankfurt am Main), 42, 329 (1964). H. S. Green, A. Maccoil, and P. J. Thomas, J. Chem. SOC., 184 (1960). C. S. Elliott and H. M. Frey, J. Chem. SOC.A , 900 (1964). (a) D. G. Retzioff, B. M. Couii, and J. Couil, J. Phys. Chem., 74, 2455 (1970); (b) G. R. Brantin and H. M. Frey, J. Chem. SOC.A , 1342 (1966). (a) M. Zupan and W. D. Walters, J. Phys. Chem., 67, 1845 (1963); (b) A. T. Cocksand H. M. Frey,J. Chem. SOC.A , 2566 (1970). (a) W. Tsanq, J . Chem. Phys., 40, 1171 (1964); (b) ibid., 41, 2487 (1964); ( c ) /b/d., 42, 1805 (1965); (d) bld., 43, 352 (1965); (e) ibid.. 44, 4283 (1966), (a) A. Lifshitz, S. H. Bauer, and E. L. Resier, Jr., J. Chem. Phys., 38, 2056 (1963); (b) P. Jeffers, J. Phys. Chem., 76, 2829 (1972); ( c ) P. M. Jeffers and W. Shaub, J. Amer. Chem. SOC., 91, 7706 (1969). D. K. Lewis and M. Keii, J. Amer. Chem. Soc., submitted for pubiication. P. Jeffers, D. Lewis, and M. Sarr, J. Phys. Chem., 77, 3037 (1973).
(14) D. M. Golden, R. K. Solly, and S. W. Benson, J. Phys. Chem., 75, 1333 (1971). (15) M. C. Lin and K. J. Laidler. Can. J. Chem., 46, 973 (1968). (16) P. J. Robinson and K. A. Hoibrook, "Unimolecular Reactions," Wiley, New York, N. Y., 1972. However, note D. L. Bunker and W. L. Hase, J. Chem. Phys., 59, 4621 (1973). (17) B. S. Rabinovitch and J. H. Current, Can. J. Chem., 40, 557 (1962). (18) (a) 8. S. Rabinovitch and M. J. Pearson, J. Chem. Phys., 41, 280 (1964); (b) M. J. Pearson, B. S. Rabinovitch, and G. 2. Whitten, J. Chem. Phys., 42, 2470 (1965); (c) F. H. Dorr and B. S. Rabinovitch, J. Phys. Chem., 69, 1973 (1965). (19) J. L. Currie, unpublished data, reported at 3rd International Symposium on Gas Kinetics, Brussels, Aug 1973. (20) I . Oref, D. Schuetzle, and B. S. Rabinovitch, J. Chem. Phys., 54, 575 (1971). (21) J. D. Rynbrandt and 8. S. Rabinovitch, J. Phys., Chem., 75, 2164 (1971). See also K. C. Kim and D. W. Setset, ibid., 77, 2021 (1973). (22) S. J. Riley and D. R. Herschbach, J. Chem. Phys., 58, 27 (1973). (23) J. M. Parsons, et a/., J. Chem. Phys., 59, 1402, 1411, 1427, 1435 (1973). (24) R. J. Gordon and M. C. Lin (US-NRL); M. J. Kurylo, et ai. (USNBS), prepublication reports. (25) S. H. Bauer, D. M. Lederman, E. L. Resier, Jr., and E. R. Fisher, lnt. J. Chem. Kinet., 5, 93 (1973).
Thermal Decomposition of Lithium Perchlorate. 1. The Initiation Rate
H. F. Cordes* and S. R. Smith Chemistry Division, Research Department, Naval Weapons Center, China Lake, Calitornia 93555 (Received September 4, 7973) Publication costs assisted by the Naval Weapons Center
The isothermal decomposition of liquid lithium perchlorate in the presence of AgC104 has been studied between 250 and 400". The final products are primarily LiCl (precipitated as AgC1) and 0 2 . Some Cl02 is produced and C103- is produced as an intermediate. The effect of added lithium chlorate was also studied. The relationship between this decomposition and that of the solid alkali perchlorates is discussed.
Introduction In a recent paper1 the authors have reported data on the initial decomposition of the solid alkali perchlorates. The initial pseudo-first-order rate constant was found to have a low Arrhenius preexponential factor. Arguments were presented to show that this low preexponential factor could be due to restriction of the rotation of the activated complex by the solid lattice.2 The data were consistent with the reaction 2c104-
-
c1*0:-*
-
2c103-
+ 0,
It was surmised that the activated complex would be freer in a liquid medium and that this reaction would then have a higher preexponential factor. On the other hand, if the low preexponential factor was associated with a "forbidden" transition in the activated complex then the rate in the liquid medium should be essentially the same as in the solid. The lowest melting of the alkali perchlorates, LiC104, was the logical choice for the test. Markowitz and Boryta3-4 have studied the decomposition of LiC104 and they found that the reaction was autocatalytic and that the autocatalysis was affected by C1-
and by AgN03. Recently, Solymosi5 has reported similar results. Neither set of investigators, however, had measured the initial rate. The present work reports on the initial rate and related phenomena. Additional data on the autocatalytic rate will be reported elsewhere.6 Experimental Section Mass Spectrometer Measurements. These measurements were made with a Bendix time-of-flight mass spectrometer and the procedure has been described in a previous paper.l All samples were in direct contact with quartz only. One important change was forced by the liquid nature of the decomposing medium. For certain rates of gas evolution there was considerable bubble formation and the gas evolution rates were very erratic. A Lectrocount integrator was connected to one of the analog outputs from the mass spectrometer. This arrangement allowed a time average (usually over 4 min) to be taken for the 32 amu peak height. Materials. The alkali perchlorates were anhydrous reagent material from the G. Frederick Smith Co. The LiC104 showed less than mole fraction of c103-. The The Journal of Physical Chemistry, Vol. 78, No. 8, 1974
774
AgC104 was hydrated reagent grade also obtained from the G. Frederick Smith Co. I t was dehydrated under vacuum and stored in the dark. The Ba(C103)~and Li2SO4 were reagent grade from Baker and Adamson. The LiC103 was prepared by mixing stoichiometric amounts of concentrated solutions of Ba(C103)~ and L i ~ S 0 4in HzO.\The C103- analysis for this material was 98.2% of theoretical, using a benzidine method.? Markowitz and Boryta4 had noted that the addition of AgNO3 to the LiC104 inhibited the autocatalysis. However, the NO3- seemed to interact with the C104-. To avoid this problem AgClO4 was used in place of AgN03 in the present work. The AgC104-AgC1 system is stable below 400O.8 At higher temperatures the AgCl melts and decomposition is accelerated. The level of AgC104 addition was between 4 and 8 mol 70.
Results Stoichiometry. The gases evolving from the decomposing LiC104 between 260 and 430" were in all cases at least 98% 0 2 , as determined by mass spectrometry. For mixtures containing AgC104 the only other gaseous product observed was ClOz. There was no sensitivity standard for C102; however, based on peak height ratios alone, the ClO2 rate was estimated a t 0.5-1.0% of the 0 2 rate. This ratio was nearly constant in time. In all cases some COz was observed. This may have amounted to 0.5-1% of the total gases, and was most pronounced at the start of a decomposition. The COz never completely disappeared from the spectrum. Some H2O was also noted at the start of a decomposition even after evacuation a t room temperature over night. The H20 was down to background after about 30 min at temperature. Small amounts of the 0 2 peak) of Li (7 amu) were seen in the vapors passing through the mass spectrometer. Samples of LiC104 containing no AgC104 were decomposed to as high as 50% decomposition and the residues quenched to room temperature. Analysis of the residues showed C1- and C103- but no C10- or ClOz-. These latter two species were analyzed for by the method in Kolthoff and Belcher,g and found to be less than anion mole fraction. The C1- and C103- were major products and the results for these materials are reported elsewhere.6 Rate Measurements. The rate measurements were made using the mass spectrometer to monitor the 0 2 evolved and are described elsewhere.1 In contrast with the work on the other alkali perchlorates, less COz was evolved during decomposition and a pretreatment a t low temperature was not found to be needed. A thermistor-controlled tube furnace was brought to temperature and then raised around the sample. The warm up period was measured to be about 10-15 min. In all cases, an arbitrary zero time was taken a t 12.5 min after the furnace was raised. Decomposition Rates for LiC104-AgC104 Alone. The rates for this system showed an initial rate that decayed in a first-order fashion to a rate that was constant in time. This behavior was apparently due to the presence of a small amount of C103- impurity (see below). After the rate had become constant in time, the sample was subjected to temperature cycling in the same fashion as for the previous samples of solid alkali perch1orates.l The Arrhenius parameters are listed in Table I. A graph is available in microfilm (see paragraph a t end of paper regarding supplementary material), No hysteresis was observed during the temperature cycling. When the temperThe Journal of Physical Chemistry, Vol. 78, No. 8. 7974
H. F.
Cordes and S. R. Smith
ature was raised above 400" the rates rose abnormally and began to accelerate with time. Experiments were discontinued when this occurred. Decomposition Rates with Added Chlorate. Chlorate is one of the products and the behavior of systems containing added LiC103 is pertinent to the discussion of mechanism. The rate of 0 2 evolution, after subtraction of Ro, the rate for LiC104 and AgC104 by themselves, was found to decay in a first-order fashion. A typical rate constant had a standard deviation of 1%.A graph is available in microfilm. If the first-order rate constant is C4 then the initial rate is found to be nearlylo dX(O,)/dt = (3/2)C4X4'
+ Ro
At the higher temperatures (>390") the first-order plots tended to curve upward at long time, indicating either decomposition of the AgC104-AgCl phase or solubility of the AgCl in the LiC104. Not all of the initial rates were as high as expected, as at higher temperatures, as much as one-third of the C103- appeared to have reacted during the warm up period. This effect was not quantitatively reproducible. The suspected cause of the problem was impurities in samples. At the lower temperatures, the values of C4 were obtained from a single sample containing 9.73 x 10-3 anion mole fraction of C103-. The rate was low enough so that over long periods of time the amount of C103- did not change. The constant C4 was given by
The amount of C103- present was large enough to that Ro was negligible as compared to dX(Oz)/dt. The run was cycled up and down in temperature and no hysteresis was observed. For this sample apparently no C103- was consumed during the warm up period as the C4 values lie on the same Arrhenius line as the other C4 values calculated from the slopes of the first-order plots. The Arrhenius parameters for the combined data on C4 are given in Table I. A graph is available in microfilm. The mixtures of LiC104-AgC104 alone behaved as if there was about 5 x 10-4 mole fraction C103- present initially. For this sample C4 = 4 x 10W4 sec-l at 344" as sec-1 from Figure 3. compared to 2.5 x Miscellaneous. Several runs were made by diluting the LiC104 with about 0.5 mole fraction of each of the other alkali perchlorates. About 0.05 mole fraction AgC104 was also present in each sample. At 380" the limiting rates were indistinguishable from that for LiC104 and AgC104 alone. These points are shown in the figures in microfilm.
Discussion of Results Chlorate Decomposition. The time dependence of the rate of 0 2 evolution from the samples containing added C103- indicate that the chlorate is disappearing in a firstorder reaction. The total amount of gas evolved over and above that expected for the LiC104-AgC104 system gives a stoichiometry of c103-
Reactions such as c103-
+ c10,-
or 2c103-
+
c1--+ 3/202
-
-
cloz- +
c102-
c103-
+ c10,-
+ 02
775
Initiation Rate of LiClO4 Decomposition
B. Monomolecular Initiation.
TABLE I: Comparison of Arrhenius Parameters for Several R a t e Constants"
Clod-
-~
Constant
Compound
A E ~
Combined M +C104 Solid only LiC104d LiC104
Ash
C4
Log A'
uAB
45.95
0.24 7.711
0.42
9.835
0.39 0.45 0.16
must occur, if at all, only to a small extent as they would drastically affect the stoichiometry. The fact that C104is produced from the decomposition of pure KC10311 would argue that the last reaction may be significant when the C103- concentration is high. A reasonable mechanism for the oxygen evolution from C103- in LiC104 in the presence of Ag+ is12
% C10-
-
+ O2
AH
+ O2
A H = -29
-2
=
followed by 2C10or
2C1-
+ C1+ 0,
2 ~ 1 0 - -+ C10,C102- -+ C1-
or
c10-
-
+
AH AH
+ A kcal/mol +A
kcal/mol
=
-5 kcal/mol
=
-24
+A
+
kcal/mol
AH = 18 kcal/mol c102c1- + 02 AH = -24 A kcal/mol In any event, a steady state for intermediates leads to C103-
-----t
2C102-
+
dX(O,)/dt = (3/2)C4X4 where C4 = K4 or 2K4 depending upon the exact mechanism. The C1- is removed from the system as AgC1. C4 cannot be a simple multiple of the rate constant for the reaction. C103-
--+
C10,-
+
0
AH
=
67
+
t
kcal/mol
as the observed activation energy for C4 is about 50 kcal. Lithium Perchlorate Decomposition with C1- Removed. Two mechanisms are presented. The first mechanism has a bimolecular initiation and the second has a monomolecular initiation. Both mechanisms were designed to be consistent with the C103- decomposition and to be as simple as possible. In all cases the C1- is removed as AgC1. A. Bimolecular Initiation. A
2C104- A 2C10,-
+ O2
A H = +14 kcal/mol
+A
This reaction is to be followed by the net result of the C103- decomposition C103-
+ c10,-
c10,-
6.909 40.10
+ O2
c10,-
ulogAC
a The Arrhenius parameters are recorded to more significant figures than appears warranted by the standard error estimates (6). This is a mathematical requirement if the best fit to the data is to be reproduced from the equations. kcal mol-' OK-*. A in sec-1. Observed initial rate divided by 4.
Cl0,-
ClO;
2 C1- + (3/2)02
C4 being either K4 or 2K4. If the mole fraction of clod- is nearly constant over the course of the oxygen evolution then one cqn derive an expression for the rate of 0 2 release. dX(O,)/dt = 4A6X5' - (3/2)(2AX$ - C,X;) exp(-C,t) For a time long enough so that exp(-Cat)
