4328
R. 34.WINGAND G. 34. HARRIS
Mechanism of the Solid-state Thermal Decomposition of Potassium Trisoxalatochromium(111) Trihydrate
by R. M. Wing1 and G. M. Harris Department of Chemistry, State Unizersity of New York at Buffalo, Buffalo 14,New Yorh (Received July 20, 1366)
The stoichiometry of the decomposition was studied by use of a simple nitrogen-swept tube furnace, with an absorption train for gas analysis; solid residues were examined by a combination of (‘met” chemical analysis, X-ray powder diffraction, and infrared spectrophotometry. The dehydration of the complex is complete a t 150°, far below the temperature required for a measurable rate of decomposition of the anhydrous material (360”). This latter process takes place, in the absence of oxygen, according to the scheme
K3Cr(C2O&
+
-
+
+
+
3 / ~ K ~ C ~ 0 l4/ ~ C r ~ 0 3 3 / ~ C O ~ 3/&0
K2C204 CrZO3----f carbonatochromium surface complex
+ KZCO3 + CO
carbonatochromium surface complex +Crz03
(1) (2)
(3)
The rates of the above reactions were determined manometrically on a vacuum furnace assembly, supplemented by periodic solid residue analysis. All obey first-order kinetics. Some carbon-13-carbon-12 discrimination studies provide strong evidence that the ratecontrolling step in reaction l is G O bond cleavage, following a short initial inhibition period. Reactions 2 and 3 occur on the subcrystalline Crz03 surface, as proven by spectroscopic identification of the surface intermediate, C02 gas pressure-dependence studies of the rate constant of ( 2 ) , and separate measurements of there action rates of pure K2C204 and hand-made Cr208-K2Cz04 mixtures. Possible structures of the surface species and its mode of formation and disintegration are discussed.
No previous kinetic studies have been made of the thermolysis of solid oxalatochromium(II1) complexes, although some significant information has been obtained by t.g.a. and d.t.a. techniques.2 Data of this latter type, however, do not in themselves offer clearcut chemical evidence concerning the nature of the reactions involved. Some recent papers by Yankwich, et ~ 1 . have ~ 3 reported complete studies of the decomposition kinetics and carbon isotope effects in the pyrolysis of simple oxalates of the type MC204 (M = Mg, Mn, and Zn), where the over-all reaction is invariably equimolar production of MO, CO, and COZ. In the present work, detailed stoichiometric and kinetic experiments were performed on racemic crystalline R3Cr(CZO4)~.3H20.The techniques used included gaseous product determination and a combination of The Journal of Physical Chemistry
“wet” chemical analysis, X-ray powder diffraction, and spectral examination of the solid residues. The results of these experiments, supplemented by some carbon-12-carbon-13 discrimination studies, lead to a reasonably definitive conception of the process. I n contrast with the “simple” oxalates mentioned above, both metal oxide and metal carbonate products are obtained, and the process occurs in several well-defined steps. One of these sub-reactions involves chromic (1) Work done by R. XI. Wing as part of the Ph.D. requirement of State University of New York at Buffalo, June 1964. Complete dissertation available from University Microfilms, Ann Arbor, Mich. (2) (a) W. W. Wendlandt, T. D. George, and K. V. Krishnamurty, J . Inorg. Nucl. Chem., 21, 69 (1961); (b) N. Dollimore and D. Nicholson, ibid., 25, 739 (1963). (3) P. E. Yankwich and P. D. Zavitsanos, J . Phys. Chem., 69, 442 (1965), and other papers referred to therein.
THERWIOLYSIS O F SOLID
OXALATOCHROMIUM(III)COMPLEXES
oxide promoted surface catalysis of oxalate ion decomposition.
Experimental Section Crystalline KBCr(C204)8-3Hz0 was obtained by standard procedure* and purified by repeated recrystallization from water. Analyses for water and oxalate confirmed its purity, which was also checked by spectrophotometric niea~urements.~When requiredl, carbon dioxide gas was obtained by sublimation of Dry Ice. ,411 other gases were taken directly from highpurity commercial cylinders. The metal oxides and graphite used were of reagent grade. The stoichiometric data were obtained by use of a simple tube furnace coupled with a conventional absorption train. A Pyr-0-Vane controller maintained the furnace within 3~0.5' of the desired temperature. The weighed samples were first introduced into an unheated zone of the furnace assembly, while the system was flushed with pure dry nitrogen, which served a,s the driver gas. To commence a run, the sample was pushed into the heated zone and its temperature was continuously recorded by means of a thermocouple attached to the Pyrex glass sample holder. The normal "heat-up', time was only a few minutes. Water determinations were made by both weight loss of the sample and collection of the escaping water vapor on anhydrous magnesium perchlorate. Carbon dioxide was determined by absorption on Ascarite and carbon monoxide similarly, after its oxidation to carbon dioxide on copper oxide wire a t 300'. The solid residues were analyzed far oxalate by permanganate titration, for chromium by iodimetry following peroxide oxidation, and for carbonate by treatment with dilute acid, collection of the evolved carbon dioxide in alkali, and back-titration. Uncomplexed oxalate was distinguished from complexed by precipitation of the former as the calcium salt, and uncomplexed chromium from complexed by use of a Dowex50X cation-exchange column. Infrared spectra of the solid residues were run in Nujol mulls on a I'erkin-Elmer Model 21 spectrophotometer, and visible and ultraviolet spectra on a Beckmann Model IIU instrument. X-Ray powder patterns of capillary-mounted samples were obtained on a Norelco X-ray diffraction setup, using either photographic or ionization-chamber recording techniques. Surface area measurements were made on a BrunauerEmmett-Teller apparatus, using nitrogen as a reference gas.6 Carbon-12-cmbon-13 isotopic discrimination data were obtained by conventional isotope ratio mass spectrometry of carbon dioxide sample^.^ Most of the kinetic r u n s were made in a vacuum
4329
furnace apparatus by following the rate of gas evolution manometrically. A magnetically controlled sample introduction device was employed, and samples of about 200-mg. size were introduced which were estimated to reach temperature equilibrium in the furance almost immediately. A liquid nitrogen cooled trap effected the separation of COz from CO, gas handling being facilitated by a Toepler pump assembly. The mercury manometer readings were made with a cathetometer and were accurate to about 0.001 cm. in a total displacement of 1 or 2 em. per run. The gas pressure data were converted to molar units by use of the ideal gas law, with corrections for the heated volume of the gas-handling system. This volume constituted about 10% of the total.
Results and Discussion ( A ) . Dehydration of the sample goes to completion in a short time near 100'. If done slowly on a thin sample below 100°, a noncrystalline product is obtained. However, if the process is carried out more rapidly by raising the temperature to 110' and a thick sample is used (or a thin sample in a sealed ampoule), a fully crystalline anhydrous salt is obtained. It thus appears that the crystallinity of the anhydrous complex is dependent upon the existence of hydrothermal conditions during dehydration. X-Ray examination showed that the anhydrous material has a completely different crystal habit from that of the hydrate (triclinic rather than monoclinic, two formula weights per unit cell rather than four). No kinetic studies were made of the dehydration process, but it clearly comprises + nucleation-crystal growth phenomenon, at least in the presence of re~iidualwater vapor. Decomposition of the anhydroue complex does not take place until the temperature exceeds 300'. (B). T h e stoichiometry of the decomposition in an oxygen-free atmosphere was determined on the basis of the following observations. (a) CO and COz are evolved in a 1:l ratio until 1.5 moles af each per mole of complex is released. Then COZ evolution ceases, but another 1.5 moles of CO is evolved for a final total of 4.5 moles of carbon oxides released per mole of trisoxalatochromium(II1) salt. (4) J. C. Bailar, Jr., and E. M. Jones, Inorg. Syn., 1, 37 (1939). (5) Peaks were observed at 5720 and 4200 A., with molar extinction coefficients of 74.9 and 97.6, respectively. A. W. Anderson and H. Sporer, J. Am. Chem. SOC.,80, 3867 (1958), report corresponding values of 5720 (76) and 4200 (97). (6) Courtesy of Dr. D. A. Cadenhead and Mr. Rumel McCallum of this laboratory. Details reported by R. McCallum, Senior Thesis, Chemistry Department, State University of New York at Buffalo, June 1963. (7) Through the kind cooperation of the mass spectrometry group, McMaster University, Hamilton, Ont., Canada.
Volume 69, Number 18 December 1966
4330
R. M. WINGd N D G. M. HARRIS
@
50% Decomposition
@
I
800
cm-'
@ 20%
K P ,
Decomposition
of
K2C20,
@
61 X Decomposition Of
@
K2Cz0,
9 5 % Decomposition of
K,C,O,
-
I
I
800
700
800
BOO
700
cm?
900
'K,Cr(C,O,),
-
I
900
91 % Decomposition
of
of K,Cr(C,OJ,
700
900
800
700
c m-1
cm-'
Figure 1. Infrared spectra of various samples.
(b) X-Ray powder patterns of the residues show clear evidence for the existence of KzCz04crystals at the point where COZevolution ceases. However, only KzC03 crystals are present after gas evolution is complete. Crz03, which must necessarily be a final product of the decomposition, does not show in the X-ray patterns, as it is apparently too finely divided. (c) Infrared spectra of residues taken at various fractions of total reaction clearly indicate the disappearance of the coordinated Cz04, which is complete by the time COz evolution ceases, and the concomitant buildup of "free" Cz042-ion (see Figure 1, curves A, B, C, and D). One then notes a decay of the Cz02peak, without equivalent buildup of the CO32- peak, though a new peak develops near 850 cm.-l as the CZ0 4 2 - ion is used up. The 850-~m.-~ peak later decays as cos2-ion appears in quantity (Figure 1, curves E, Table I: Analytical D a t a for Partial
F, G, and H). Carbonato complexes of other transition metals are known to exhibit characteristic inso we interpret frared absorption peaks* near 850 our observations to indicate the appearance of a carbonatochromium surface complex ~ t s a reaction intermediate in the conversion of CZOd2-ion to COs2ion. (d) A comparison was made of quantities of carboncontaining products in four partially reacted samples, as calculated from the gas evolution datag and similar results based on the results of "wet" analysis for carbonate (by acid decomposition) and total oxalate (by permangante titration). The data are recorded in terms of moles of the various carbon-containing compounds present in the product mixture per mole of carbon in the original sample of trisoxalatochromium(111) complex, as shown in Table I. These data confirm the conclusion drawn earlier that some of the carbonate does appear as a carbonato complex which later decomposes. The discrepancies
Trisoxalatochromium(111) Decomposition Gas analysis dataQ C' A' B'
0.00 0.03
0.00 0.00
0.33 0.73 0.84
0.87
0.67 0.24 0.16 0.13
~
Wet analysis dataD AI/ BI/ Cft
0.05 0.30 0.17 0.23
' Fractions A, B,
0.31 0.47 0.66 0.64
0.64 0.23 0.17 0.13
Differencesb AA AE
0.05 0.27 0.17 0.23
0.02 0.26 0.18 0.23
and C refer t o the mole fraction of carbon appearing as COa2- ion, CZOP ion, and C Z O complex, ~ respecAA = (A// - AI); AB = (B' tively.
The Journal of Physical Chemistry
~~
(8) J. Figata, A. E. Martell, and K. Nakamoto, J. Chem. Phys., 36, 339 (l962), and other references given therein. (9) This calculation is based on the assumption that the total stoichiometry for complete decomposition is given by eq. 1, plus the reaction KzCz04 +. KzCOs CO. Then, if [KsCr(C~04)s]0 is the initial number of moles of complex compound taken, after partial decomposition the following relations hold
+
-
[Kt,Cr(CzO&1 = [KsCr(CzO&lo 2/~[CO~l [KzCz041 = 2[C01] [CO]
-
[CrzOsl = '/dCOzl [KzCOsl = [CO] - [Cod
4331
THERMOLYSIS OF SOLIDOXALATOCHROMIUM(III) COMPLEXES
in the carbonate determinations (AA) balance out, within experimental error, those in the oxalate ion determinations (AB), showing that the carbonato complex “lost” in the gas evolution data computation appears, as it must, as ionic oxalate. It is also obvious that no intermediate is formed in the conversion of complexed oxalate to “free” oxalate ion since AC is zero within experimental error. (e) Finally, it was confirmed by “wet” analysis that completely reacted residues contain 1.5 moles of C032- per mole of chromium(II1). This, when added to the total of gaseous carbon oxides evolved, completes the carbon balance for the system. A suitable reaction Bcheme to interpret the material balance data for decomposition in a nitrogen atmosphere follows. Decomposition of Anhgdrous Trisoxalatochromium( I I I )Complex
+
K3Cr(C204)3 +a/~Md3204
1/&r20a
+ a/2C02 + 8/2C0
0.2
0
40
20
IO
TIME,
50
60
MIN.
Figure 2. Typical first-order kinetic plot for the complex decomposition.
(1) IO~T-‘
Catalyzed Potassium Oxalate Decomposition
1.50
1.48
1.46
+
1.52
1.56
1.54
1.58
K ~ C Z O ~C r ~ 0 3+ carbonatochromium surface complex (2) carbonatochromium surface complex + Crz03
+ KzCO~+ CO
(3)
A few decomposition experiments were carried out ;in air instead of nitrogen. Invariably, only 0.5 mole of COs2- remains at complete reaction per mole of chromium, and the latter all appears as chromate. Clearly, “carbonate fusion oxidation” occurs under these conditions, according to the over-all stoichiometry 2K~Cr(C20.&
+ 3/20#2 -+
KG03 4.2KzCr04
N
2
2.0
-
L
0 c
0.7 1.0
+ 6CO + 5co2
0.9 0.8 -
(4)
No study has been made of the kinet,ics of this process in the present work. (C). T h e kinetics of the trisoxalatochromium(II1) complex decomposition, (reaction 1) is clearly first order after a brief induction period. As indicated by the infrared spectra and analytical data discussed above, there is no evidence for intermediate products of any kind. A plot of the data for a typical run is presented in Figure 2. The first-order rate constant was found to be only slightly dependent on the addition of inert gases to the system, and equally so for argon, CO, and COZ. For example, a, pressure of 48 cm. of initially added COZreduced the rate constant from the in uacuo value of 5.0 X 10+ mjn.-l ah 400’ to 3.5 X 10-2mh-l.
i
:,i0.6
0.5 0.4
-
0.3 1.50
1.52
1.54
1.56
1.58
1.60
IOaT’’
Figure 3. Temperature dependence of the induction period, bind, and the rate constant, B, for the decomposition of anhydrous K&r( CzO&.
Both the induction period and the decomposition rate constant are strongly temperature dependent (see Figure 3). The induction period has an Arrhenius activation energy of 59 kcal./mole. The EyringPolanyi activation parameters for the decomposition Volume 69,Number 18 December 1966
4332
R.
Table I1 Species
co
cos
KaCaOi
C18 atom % excessa ( x 106)
-15 1 1
8*1
5 1 1
a As compared to a standard Cog sample with a carbon-13 atom percentage of 1.113.
reaction are AHI* = 34 kcal./mole and A&* = -21 e.u. The isotope e$ect study consisted of determining the carbon-12-carbon-13 ratios in each of the carbon-containing products of reaction 1. The averages of three independent sets of data, taken from experiments in which the complex oxalate was 50% decomposed, are shown in Table 11. It must in logic be assumed that each pair of CO and C02 molecules produced in reaction 1 originates from a single C204unit. These units must therefore, on the average, have an isotopic composition of about -4 X atom % excess of carbon-13, which comes close t o balancing outlo the value for the undissociated ( 2 2 0 4 observed 5 X units appearing in the K2C204. It seems likely that the initial step of the reaction which produces the oxides of carbon is oxygen-carbon bond cleavage, with the 0-Cl2 bond favored over the Q4Y3 in the conventional manner.ll Succeeding steps then produce CO and COz, with the CO always originating from the end of the oxalate group which was the site of the initial 0-C bond cleavage. The total mechanism of the trisoxalato complex ion decomposition is more difficult to visualize. There are two distinct possibilities which lead to the observed product distribution. One comprises chromium-oxygen bond breakage12 a t one end of a chelated oxalate group, followed by release of the carbon oxides and rearrangement of the amorphous oxalate ion-chromic oxide-potassium ion residue into crystalline K2C204 and subcrystalline Cr203. Such a mechanism is supported by data on the aquation of Cr(C204)3a-ion in acidic s01ution.I~ Since the reaction takes place in aqueous medium, the coordination positions reversibly vacated by one-ended dissociation of oxalate groups are immediately re-occupied by water molecules. This preserves octahedral symmetry in the activated state for complete release of oxalate ion, thus causing little change in the crystal field stabilization energy, since water and oxalate have very nearly identical crystal field strengths. I n the solid-state decomposition reaction, however, a square-pyramidal activated State is almost a necessity, resulting in a lOSS Of CryStal The Journal of Physical Chemistry
n!r.
WING AKD
G. 31. HARRIS
field stabilization energy of about 10 kcal./mole.14 It is interesting that the activation enthalpy of the thermal decomposition does exceed that of aquation by about this amount (34 kcal./mole as compared to 23 kcal./mole). Since it appears possible that oxalate ions are released intact from the decomposing trisoxalatochromium(II1) complex, a mechanism can be suggested for the thermolysis in purely ionic terms. It can be visualized that half of the oxalate ions, on release from the complex, break up by an ion-radical mechanism to yield CO, GOz, and 0 2 - ions, while the remainder pair up with K f ions to yield crystalline K2CZO4. “Free” oxalate ions with the two electronic charges divided equally between the four oxygens would have a carbon-carbon bond energy of less than 10 kcal./ mole.l5 This is much smaller than the minimum decomposition enthalpy defined by the observed enthalpy of activation (34 kcal./mole), so it is obvious that the positive ions must play a critical role in stabilizing the oxalate ions which survive as K2C204. T h e catalyxed potassium oxalate decomposition, (0). symbolized by reactions 2 and 3, was studied by using samples of trisoxalato complex which had been heated until all complexed oxalate was decomposed according to reaction 1 ( i e . , no more C02 was being evolved). Then reaction 2 was followed by observing K2C20( decomposition (determined by permanganate titration) and reaction 3 by manometry of carbon monoxide evolved. The data were treated on the basis of successive first-order reactions. The Eyring-Polanyi activation parameters for the two processes, obtained from (10) The carbon-13 balance should be reckoned as follows, assuming no isotopic discrimination in the residual complex
+
2 A c ? 1 3 ~ 2 u s ~AC’3co r
+ AC1%os = 0
The actual summation gives 3 ( 1 2 ) X atom ’% excess, which is not quite zero. The discrepancy may result from the assumption made concerning the residual complex isotopic composition. (11) The expected isotope effect would be of the order of magnitude of 2% (see L. C. Melander, “Isotope Effects on Reaction Rates,” Ronald Press Go., New York, N. Y., 1960). The apparent value in the present study is (AC1*coe AC%o)/C1%tandard= 2.1%. (12) There is no direct evidence for this type of mechanism in the case of chromium(II1)-carboxylic chelates. However, studies of the acid hydrolysis of chelated carbonate in Co(NHa)4C03+ ion clearly support the concept of metal-oxygen bond opening in the dechelation step prior to COa liberation (see F. A. Posey and H. Taube, J . Am. Chem. Soc., 75,4099 (1963)). (13) K. V. Krishnamurty and G. M. Harris, J. Phys. Chem., 64, 346 (1960); H. Relm and G. M. Harris, to be published. (14) F. Basolo and R. G . Pearson, “Mechanisms of Inorganic Reactions,” John Wiley and Sons, Inc., New York, N. Y., 1958, p. 108 ff. (16) This estimate is based on a simple electrostatic repulsion calculation, assuming a charge separation of 5 A. (a reasonable maximum, since the diagonal 0-0 distance in oxalate ions is less than 4 A.) and a “normal” carbon-carbon bond energy of 80 kcal./mole.
-
4333
THERMOLPBIB OF Sor,xr,OXALATOCRROMIUM(III) COMPLEXES
io3T-'
15
I
-
10 9-
-
'a .-
E
8 -
7-
m "
Y N
5-
L
4-
0
0
0
4
e
6
12
10
14
IS
19
20
P
N
m";
2
0
3
-
Figure 5. Dependence of KzCz04 decomposition rate (catalyzed by CnOg) on external gas pressure.
2 -
1.5
-
io3T-'
0
Figure 4. Temperature dependence of t h e rate oon&ant for K2C204 disappearance, kz, and for COBevolution, kl, during CrzO,-catalyzed KzC204decomposition.
10
20
30
,40
'SO
60
TIME, MIN.
temperature dependence data (see Figure 4), are AHz* = 35 kcal./mole, ASz* = -25 emu.,AH3* = 46 kcal./mole, and ASs* = -9 e.u. The effect of COz pressure on reaction 2 at 420' was studied, yielding the results Pcoz, cm. 10zlcz,min.-I
0.0 4.3
2.6 3.0
6.0 1.7
20.0 1.2
An inhibition is seen to occur which is interpreted in terms of successively greater contamination of the active CrzOs surface by adsorbed COz as the pressure of the latter is increased. This interpretation is supported by treatment of the data on the basis of the Langmuir isotherm, which leads to a rate expression of the form16
p/(ko - IC)
=
a/a
+ p/ko
(5)
In this equation, p ia the pressure of COz, ko = ICz = rate constant of reaction in vacuo (as regularly determined above), and IC is the rate constant of the inhibited reaction. Figure 5 shows the application of this relation to the daka given. It is seen to be a very satisfactory confirmation of the concept, since the observed slope of the curve (26 min.) is very close to the expected value of l/k2 = 23 min. The nature of the catalysis in this process was examined by another procedure. Hand-mixed samples of finely divided CrzOa and KzC204were prepared by
TIME, MIN.
Figure 6. Kinetic plots for decomposition of a hand mixture of KZCZOA and CrzO,.
thorough grinding with mortar and pestle. The progress of the reaction of this material, determined at 450' by the usual oxalate decomposition procedure, is illustrated in Figure 6. It is noted that the catalyzed reaction is incomplete (curve A), but that the initial rapid process is first order (curve B), with a halftime of 5.7 min. The corresponding rate constant is
+
(16) The Langmuir isotherm is usually written (1 - 8) = u / ( a p ) , where 0 is the fraction of surface covered, a is a constant determined by the relative rates of adsorption and desorption, and p is the pressure of the gas being adsorbed. If it is assumed that the rate constant of reaction on the surface, k, is proportional to the fraction of surface available, then k = ko(1 e), where ko is the maximum rate constant, achieved in vacuo. Substituting for (1 - 8) and rearranging, one derives eq. 6.
-
Volum 69,Number 1% December 1066
4334
R. 41. WINGAND G. M. HARRIS
C
4‘0
A6
r,
0
160:- t C O
tc0:-
Figure 7. Mechanisms of reaction.
0.12 min.-l, identical with the kz value at this temperature for Cr~O~-K&~04 mixtures prepared in situ by
K&r(C2Od)s decomposition. The only difference is the lack of completeness of the reaction of the handmade mixture, undoubtedly because of the impossibility of achieving totally intimate contact between Cr203surfaces and C A 2 - ions by this procedure, The ir, situ material, however, decomposes completely by the catalymd mechanism, since here total mixing has been achieved by the preparation procedure, as shown by the extremely high state of flubdivision of the material.17 Good first-order kinetics are thus observed over several half-times of reaction of the in situ preparations. The effectiveness of some other substances as catalysts for KZCzO4 thermal decomposition was also studied by use of the handmade mixture technique, The results of the experiments, all performed a t 480”, were
The metal oxides or@all notably better catalysts than graphite, This supports our concept of a metdcarboneto surface complex intermediate, formed by addition of C0,- radicals to the electron-accepting metal oxide surfaces at exposed oxygen sites. Also, the catalytic activity of the three oxides falls into an order similar to their n-type character order of Zn0 > Cr208 > AI2O3. Graphite can act as an electron acceptor, but can hold the COz only as adsorbed gas, not as a carbonato complex species, and so is less effective by one order of magnitude. One can visualize the completion of the catalysis process in either of two ways, both of which assume initial splitting of CZOd2- ion into C02- radical ions, with simultaneous attachment of the latter as carbonato
The Journal of Physical Chemistry
groups on adjacent surface oxygen sites. In one mechanism, Goaz- and CO are liberated by simple exchange of oxygen between the radicals (Figure 7%). I n the other, an oxide vacancy must be temporarily formed in the metal oxide surface (Figure 7b), A choice between these mechanisms might be possible by means of experiments with oxygen-18-labeled oxalate, though other paths of oxygen exchange might obscure the result, Carbonato surface complexes have been proposed before, in connection with the oxidation of CO to C02 on metal oxide ~ u r f a c e d including ,~~ Some supporting infrared data. 2o However, Courtois and Teichner2I do not concur in the carbonato complex concept for the CO 0 2 reaction on nickel oxide catalyst. They may well be correct in reference to carbon dioxide adsorption, since it has linear geometry. However, the COBradical ion of our concept mtiofies the geometrical requirements much better, since it has the saine C,, point symmetry as a carbonate groups8# Alao, the excess electron should promote strong bond formation with surface oxygen, Finally, some comment should be made concerning the significance of the first-order kinetics observed in all three reactions atudied. Uwally, this type of solid-state kinetics is associated with random nucleation by single-molecule decomposition as the rate-controlling This is fully consistent with our identification of the rate-controlling steps in all three reactions as intramolecular processes involving only one ionic species a t a time, following an equilibration period of buildup of decomposition nuclei. The latter process accounts for the Ahort inhibition time observed in the study of reaction 1.
+
Acknowledgment, Financial support of this work through Contract No. AT(30-1)-1578 with the U, S , Atomic Energy Commission is gratefully acknowledged, Thia paper constitutes Report No. NYO-1578-32 to the Commission. (17) The surface area of these preparations was between 10 and 16 m.Q/g., as determined by B.E.T. procedure.6 (18) An extrapolation of the data of S. Akalan, Em. Fac. Sci. Univ. Istanbul, 21C, 184 (1966), leads to a value of k = 1.4 X 10-4 m h - 1 , in good agreement with our determination at 450’. (19) See F. 8 . Stone in “Chemistry of the Solid State,” W. E. Garner, Ed., Butterworth and Co. Ltd., London, 1966, Chapter 15, p. 397 ff. (20) R. P. Eischens and W. A. Plisken, Adwan. Catalysis, 9, 662 (1967). (21) M. Courtois and S. J. Teichner, J . CataZysis, 1, 121 (1962). (22) D. W. Ovenall and D. H. Whiffen, Mol. Phgs., 4, 136 (1961). (23) See ref. 19, Chapter 7.
