Kinetics and Mechanism of the Thermal Decomposition of Hexaamminecobalt(lI1) and Aquopentaamminecobalt(II1) Ions in Acidic Aqueous Solution

ANTHONY MARTIN NEWTON AND THOMAS WILSON SWADDLE' Depurtnlent nf Chernirtry, The U t z i ~ ~ r s i t y of Calgnry, Cnlgcrr~ , Alhertn T2N IN4

Received March 18, 1974

ANTHONY MARTIPI. NEWTON and THOMAS WILSOU SWADDLE. Can. J. Chern. 52,2751 (1974). The initial step in the thermal decomposition of CO(NH, )?~+ in acidic aqueous solution is the

replacement of NH3 by H 2 0 , which occurs by a hydrogen-]on independent path, first order in complex, with rate coefficient k , = 7.9 x lo-' s-' (140.4"), AH* = 36.6 kcal mol-', and AS* = 10.7 cal deg-' mol-' in 0.1 M HCIO,. For CO(NH,),OH,~+, there is a similar initial aquation path with k l = 12.6 x lo-' s-' (140.6'), AH* = 41.9 kcal mol-', and AS* = 24 cal deg-' mol-' and also a path first order in complex but inverse first order in [H+] with k2, = 6.2 x lo-' M s-' (140.6'), AH* = 43.5 kcal mol-', and AS* = 26.7 cal deg-I mol-I, in per- chlorate media of ionic strength 1.0 M. The effects of electrolyte type and concentration on the rates of these reactions have been examined. Subsequent aquation steps are relatively rapid because of the predominance of inversely [Hf]-dependent pathways and are followed by redox to CO(H,O),~+, NH,+, N2, N 2 0 , and a minor amount of 0 2 . A mechanism involving OH and NH2 radicals is proposed for the redox step.

ANTHONY MARTIN NEWTON et THOMAS WILSON SWADDLE. Can. J. Chem. 52,2751 (1974). L'etape initiale lors de la decon~position thermique du CO(NH,),~+ en solution aqueuse

acidifite est le remplacement d'un NH, par une molecule d'eau; cette reaction se produit par une voie n'inpliquant pas d'ions hydroghe, elle est du premier ordre en complexe avec un coefficient de vitesse k , = 7.9 x lo- ' s - ' (140.4"), AH* = 36.6 kcal mol-I et AS* = 10.7 cal deg-' mol-' dans HCIO, 0.1 M. Dans le cas du CO(NH,),OH,~+, il existe une equation similaire pour le chemin initiale et k, = 12.6 x s-' (140.6"), AH* = 41.9 kcal mol-' et AS* = 24 cal deg-' mol-'; cette reaction est aussi du premier ordre en complexe mais d'un ordre inverse du premier en [H+]avec k2 ' = 6.2 x lo-' M s - I (140.6"), AH* = 43.5 kcal mol-' et AS* = 26.7 cal deg-I mol-' dans un milieu perchlorate ayant une force ionique de 1.0 M. On a aussi examine les effets de type tlectrolyte et de concentration sur les vitesses de ces reactions, les tquations des Ctapes subsequentes sont relativement rapides a cause de la predominance d'un chemin dependant d'une f a ~ o n inversement proportionnelle a la concentra- tion [H+ 1, ces Btapes sont suivies par une reaction d'oxydo-reduction conduisant a C O ( H ~ O ) ~ ~ + , NH,+, N2, N 2 0 et des quantitks mineures de 0,: On propose un mecanisme impliquant des radicaux OH et NH, pour I'ttape d'oxydo-reduction. [Traduit par le journal]

Introduction A striking feature of cobalt(II1) chemistry is

the great decrease in reactivity which occurs on replacing several of the aquo ligands in Co- (HzO)63+ by NH,, to the extent that the ammine ligands in CO(NH,),~+ and CO(NH,),X(~-")' (X = halogen, H,O, oxyanions, etc.) are com- monly regarded as substitution inert in acidic aqueous solution. Nevertheless, kinetic studies of the aquations of X n - from CO(NH,),X(~-")' begin to suffer from the complicating effects of NH, loss from either the parent complex or (more usually) from the product Co(NH,),- OH,3+ at temperatures above 80" (1). The present study seeks to determine the importance

'To whom correspondence should be addressed.

of this complication, as well as to gain insight into the mechanism of decomposition of Co- (NH,),,' and Co(NH3),0Hz3 + in solution.

Since the completion of our kinetic studies, Garner and his co-workers (2) have published data on the decomposition of Co(NH3),OHZ3+ and ~~S-CO(NH,) ,(OH,)~~ + in acidic perchlorate media, in addition to their papers on the decom- position of the presumed facial isomer of CO(NH,),(OH,),~+ (3) and cis-Co(NH,),-

+ (4). It appears that the initial replace- ment of an ammine ligand by water is the rate- determining step in the decompositions of the penta- and tetra-ammines (2) but that redox decomposition to cobalt(I1) predominates in the decomposition of the triammine (3) (however, up to 17% of the observed rate of disappearance

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2752 C A N . J . C H E M . VOL. 5 2 , 1974

of the tetraammine could also conceivablv be due to redox (2)). cis-Diamminetetraaquocobalt- (111) decomposes by direct redox to cobalt(I1) (4), whereas amminepentaaquocobalt(I1I) under- goes an extraordinary disproportionation to cobalt(I1) and diamminetetraaquocobalt(II1) (5).

We shall address ourselves primarily to two salient questions which remain. Firstly, there is no information on the kinetics of decomposition of the hexaammine, and, secondly, the nature of the oxidized reaction products has not been established. The latter has an important bearing on the mechanism of the redox steps. In addition, new information pertaining to electrolyte effects on the reaction rates will be presented, together with improved activation parameters for the decomposition of CO(NH,),OH,~+.

Experimental iClaferial~

Salts of lithium, CO(NH,),~'. and Co(NH3),0HZ3+ were made by standard methods (6, 7) and were checked for purity by chemical analysis and by examination of their visible spectra; these and all other spectral measure- ments were made using a Cary Model 15 spectrophotom- eter. Baker Analyzed perchloric acid (72%) and Fisher "purified" sodium perchlorate monohydrate and sodium nitrate were used directly. Distilled water was either passed through Barnstead deionizer and organic removal cartridges or else redistilled carefully from alkaline per- manganate before use; the kinetic results were the same in either case.

Kinetic Studiies Aliquots of solutions of the appropriate complex in

aqueous NaC10,-HC10, or LiCI0,-HC10, of the re- quired ionic strength I were sealed into Pyrex ampoules, and these were preheated to about 95 (to facilitate ther- mal equilibration and to dissolve any solid hexaamn~ine- cobalt(II1) perchlorate) before immersion in an oil-filled Lauda NS-HT thermostat bath (-t0.lC). Light levels in the bath were negligible. Timing of the reactions was begun 2 min after inlnlersion of the samples, and ampoules were withdrawn periodically and chilled to room temperature. The samples were then analyzed spectrophotometrically at 427 nm (for the hexaammine) or 344 or 490 nm (for the aquopentaammine). Alternatively, the cobalt(I1) content of the samples were determined by making r. ml of the aliquot and (6 - c ) ml 0.1 M HCIO, up to 25 ml with concentrated HCI, and measuring the optical absorbance of the resulting blue solution at 690 nm (E 471 M - ' cm-' according to a calibration curve established using solu- tions of pure cobalt(11) nitrate) (8). The cobalt(1I) con- centrations so measured were also checked in some cases by Kitsen's method (9), with essentially identical results.

Mass Spectra o f t he Gaseorts Decotnposition Prodrict~ Solution aliquots, made up as for the kinetic experi-

ments, were placed in breakseal ampoules and thoroughly degassed on a vacuum line by repeated freeze-thaw cycles.

The an~poules were then sealed and immersed in the thermostat bath at 130.6' for several half-periods of the decomposition reaction. The samples were then frozen at liquid nitrogen temperatures and the gaseous contents of the tubes were released into a VarianIMAT CH-5 mass spectrometer for analysis.

Results Throughout this article, the molar concentra-

tions cited refer to solutions as at 25". Uncer- tainty limits quoted represent standard errors; where these are not stated for an experimental measurement, a single determination is implied.

Decor?zposition of Aquoperztaamminecobalt(III) Ion

The spectrum of the final reaction products (absorption maximum at 505 nm, E - 5 M-I cm-l) identified the Co-containing species as being entirely hexaaquocobalt(I1) (10). The absorbance A , at time t changed in accordance with first order kinetics as the reaction pro- ceeded, except during a short initial "induction period" (Fig. I ) ; the latter phenomenon was much less marked than that reported (1 1) for the deconlposition of cobalt(II1) ammines in molten NH4HS04 and in 9 7 z sulfuric acid. Further- more, as Garner and co-workers have observed also (2) , the visible spectra of partially reacted solutions showed isosbestic points at the early stages of the reaction, but these were not main- tained as the reaction proceeded. These facts ind~cate the presence of a fairly long-lived reac- tion intermediate in the decomposition.

Because CO(NH,),(OH,),~+ is known (3) to decompose rapidly under the conditions of these experiments, the intermediate was presumably a mixture of cis- and trans-CO(NH,),(OH,),~',

MINUTES

FIG. 1. T ~ m e dependence of the opt~cal absorbance A, (490 nm, 10 mm opt~cal path, 25 )of a solut~on ~nrtially 0.0106 M aquopentaamn~inecobalt(lIl) perchlorate In HC10,-NaC104 (I = 1 .O M) at 140.6'.

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NEWTON AND SWADDLE: COBALT(I1I) AMMINES

TABLE 1. First-order rate coefficients k,,, for the decomposition of C O ( N H ~ ) ~ O H ~ ~ + in 0.1 M HCIO,"

Temperature 102[NaC104] lo2 [NaN03] l ~ ~ k , , , rc) ( M ) (M) (s - I)

"[CO] ,,,,, = 0.0106 M. bComplex present as its nitrate salt; elsewhere, as the perchlorate.

TABLE 2. Hydrogen Ion concentration dependence of the first-order rate coefficient k,,, for the decon~pos~t~on of CO(NH,)~OH,~+ "

-- -- - -

Temperature [HCIO,] IO4k0b, lo5kl 105kz ("C) (MI w l ) (ss l ) (M s-')

= 1.06 M, adjusted with NaCIO,. Initial [Co(NH3),0H13-I = 0.0106 M.

with the former predominating, as Garner and K, CO-workers suggested (2). These decompose [2] C0(NH3)50Hz3+ * C0(NH3)50H2+ + H + more rapidly than their progenitor (2), and so the slopes of the linear portions of the plots of

k2/ + H.0 H +

log (A, - A,) against t gave the first-order rate C O ( N H ~ ) , ( O H ~ ) ~ ~ + + NH3 -+ NH4+ coefficient kobs for the rate-controlling reaction

H + [3] kobs = k1 + k,~, [H+l- ' = kl [I] C O ( N H ~ ) ~ O H ~ ~ + 3 C O ( N H ~ ) , ( O H ~ ) ~ ~ + + NH4+ + k , , [ H + ] - '

The values of kob, collected in Tables 1 and 2 In order to facilitate comparison with the confirm that parallel acid-independent (rate CO(NH,),~+ system (in which the low solubility coefficient k,) and inversely acid-dependent (k,) of the perchlorate salt imposes some restrictions paths operate (2), as per eqs. 2 and 3, and that on the kinetic experiments), the effect of nitrate high concentrations of perchlorate ion have no ion on the rate was also examined, and this drastic effect on the reaction rate. anion was found to cause a modest acceleration

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2754 CAN. .I. CHEM. V O L . 5 2 , 1974

of the decomposition reaction (Table 1). We attribute this to the fact that nitrate ion has appreciable Br~nsted basicity in aqueous solu- tions at high temperatures (the pK, of aqueous nitric acid rises from - 2 at 0" to + 2 at 300" (12)) and will therefore decrease the free hydrogen ion concentration and so increase the contribution of the inversely hydrogen-ion dependent path- way.

The reaction product Co2+ has been reported to catalyze the radiolytic decomposition of hexaammine- and aquopentaammine-cobalt(111) (13), and to inhibit the reduction of Co(NH,),- OHZ3' by S,0a2-/Ag+ (14). The effect of added cobalt(I1) perchlorate on the rate of the spon- taneous decon~position of CO(NH,),OH,~' in 0.1 M HC10, was therefore studied with refer- ence to control experiments in which Zn2+ was added in place of Co2+, as these ions have closely similar ionic radii and therefore similar medium effects. At 130.6", with an initial aquo- pentaamminecobalt(I11) perchlorate concentra- tion of 0.0106 M, and [M(CIO,),] = 0.0109 M, k,,, was (3.52 + 0.13) x lo-, and (3.49 f 0.06) x lo-" s-' for M = Zn and Co, respectively; the corresponding data for [M(CIO,),] = 0.109 M were (2.57 i 0.02) x lo-" and (2.61 k 0.03) x lo-" s-I, as against (4.01 + 0.01) x lo-" in

the absence of added M(I1) perchlorates. Thus, divalent metal perchlorates exert a small retard- ing effect on the reaction (as does NaC10,). but there is no specific effect attributable to cobalt(11). Nevertheless, at high concentrations of added cobalt(II), an unidentified black precipitate formed as the reaction proceeded; this was not observed when Co(I1) perchlorate alone was heated in 0.1 M perchlor~c acid at 130" for several days, which suggests that Co(I1) in high concen- trat~ons can react with one of the decomposition products of CO(NH,),OH,~+ formed after the rate-controlling steps.

Decomposition of Hexaamniinecobalt (111) Ion For this ion, as for the pentaammine, the final

cobalt-containing product was CO(H,O),~+ but the spectral changes occurring during the decom- position reaction (Fig. 2) make it clear that an intermediate Co(111) complex was again in- volved, and this is verified by the appearance of the semilog kinetic plots (Fig. 3, cf. Fig. 1). Ion exchange chromatography of a solution of partially-decomposed CO(NH,),~ + on Dowex 50W-X4 resin with 1.0 M HC10, produced a

0 440 480 520 560

WAVELENGTH nm

FIG. 2 . Spectrum (25 ' , optical path 10 mm) of a solu- tion originally 0.0106 M CO(NH,),~+ in 0.1 M HCIO,, (a) initially, and after (b) 100, (c) 360, and ( d ) 830 min at 130.3".

MINUTES

FIG. 3. Decomposition of Co(NH,),,+ in HCI0,- NaC10, ( I = 1.0 M ) at 149.6-, followed by spectro- photometric determination of cobalt(I1) by Kitsen's method (9).

rapidly-moving red band (hexaaquocobalt(I1)) and a thin, slow-moving pink band closely fol- lowed by an extensive yellow band of Co- (NH3),3t; the pink band was so small as to render its isolation impracticable, but it almost certainly consisted of Co(NH,) ,OH2, +, as the more highly aquated ammines would not have survived in detectable quantities under the experimental conditions (2, 3).

The rate coefficients k,,, (Table 3) for the decomposition of hexaamminecobalt(III) were obtained from the linear portions of semilog

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NEWTON AND SWADDLE: COBALT(II1) AMMINES

TABLE 3. First-order rate coefficients k,,, for the decomposition of Co(NH,),j+ in acidic aqueous solution

Ionic Temperature [HCI04] strength Supporting 1O4kObs

("C) (MI (MI electrolyte (S -

0.033 1 . O NaClO, NaCIO, LiCIO,

0.100 1 . O NaClO, NaCIO, LiCIO,

1 .oo 1 .o

.'[Co'+] measured by HCI method. b[CoZ+] measured by Kitsen's method (9). <Direct spectropbotometric determination. Co(h'H3I63+ present as the nitrate salt (0.0106 M ) :

elsewhere, as the perchlorate.

plots such as Fig. 3 ; this was a valid procedure, since Co(NH3),0Hz3' had been shown to de- compose some 20 times faster than its progenitor under the experimental conditions. The effect of [H'] on the reaction rate in 1.0 M perchlorate media was investigated by analyzing for the product Co(I1) rather than by attempting to mea- sure the absorbance of the surviving hexaam- minecobalt(II1) ion, since the perchlorate salt of the latter is insufficiently soluble in these media at room temperature; however, it was established at low perchlorate concentrations that both methods gave the same results.

The data of Table 3 show that the ionic strength effect was independent of the choice of NaClO, or LiCIO, as the supporting electrolyte and that the hydrogen ion concentration had only a very small effect on the reaction rate. The latter observation contrasts sharply with the case of CO(NH,),OH,~+ but is scarcely sur- prising, since the pKa of CO(NH,),~+ is about 16 at 25" (15) as against 6.5 for CO(NH,),OH,~+ (16), so that the pathway corresponding to k, (eq. 2) is inaccessible for the hexaammine. Hy- drogen ion effects have therefore been ignored, i.e., kObs has been set equal to k , of eq. 3, in com- puting the activation parameters for the hexa- ammine complex (Table 4). The kinetic param- eters obtained for the decomposition of the aquopentaammine ion are also listed in Table 4 and are in satisfactory agreement with (and

somewhat more precise than) those obtained by Garner and co-workers (2).

Gaseous Reaction Products The gaseous reaction products of the thermal

decompositions of 0.01 M CO(NH,),~+, Co- (NH3),0H,3+, and hydroxylammonium sulfate at 130.6" in 0.1 M HCIO, were identified mass spectrometrically as N, (mle 28), N,O (mle 44) and, in the case of the complexes, a minor amount of 0, (mle 32). The peak heights R, relative to that for N, being 100, were measured at an electron energy of 70 V. For hexaammine- cobalt(II1) nitrate, R = 7 and 69 for 0, and N,O, respectively ; for aquopentaamminecobalt- (111), the corresponding R values were 9 and 35 (for the nitrate salt) and 10 and 51 (for the per- chlorate), whereas for (NH,OH),H,SO, they were 0 and 29. For a sample of pure N,O, R at 70 eV is 990 for mass 44; thus, virtually all the N,' and 0, + detected mass spectrometrically were derived from the N, and 0, in the reaction products and not from the N 2 0 f . Furthermore, it was established spectrophotometrically that nitrate ion does not undergo significant decom- position in 0.1 M HCIO, at 130°, and in any event it is clear from the mass spectrometric data that the presence of nitrate in place of perchlo- rate does not affect the relative yields of N, and N,O significantly. It is therefore safe to conclude that the ultimate decomposition products of both

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TABLE 4. Klnet~c parameters for the rate-determ~n~ng steps In the decomposlt~on of cobalt(II1) ammines In 0.1 M HCIO, at Ionic strength 1.0 M

A&,* o Nature of rate controlling 106kl, . 5 r 1 A H 1 * AS1* (cal kz., M A - ' AH2 * (cal deg-I >

Complex step (109.8 ) (kcal mol-I) deg-I mol-I) (109.8") (kcal mol-I) mol-') - - - -- -- - - --

Ref. - ._- L1

C O ( N H ~ ) ~ ~ + a 2.25" 36 6 k 1 . 8 10.72 5 .4 This work 2 E

C O ( N H ~ ) ~ O H , ~ + Aquation to 1.71 41 .9k1 .3 24 4 8.1 x 43.521.8 26.7k 5.4 This work CoiNH3)4(OHz)Z3 + 2.05 37 .9k6 .0 14 & I 5 8 . 1 ~ 40 k 5 17 k 1 6 C

2 0 c~s-CO(NH,),(OH,),~+ Aquation to -1.1

CO(NH,),(OH,)~~+

~ ~ c - C O ( N H , ) , ( O H , ) ~ ~ + Redox to Co(I1) - e

c~s-CO(NH~),(OH,),~+ Redox to Co(l1) -- -- -- -

4 . 8 ~ 1 0 - ~ ~ 36 .7k0 .9 26 .12 2.7 - --

4

"Ionic 3trength 0 16 M, nltrate salt of complex used. bCalculated for 109 8" from the actwatton parameters, for cornparlaon wlth d ~ r e c t nicasurements made by Garner rt a/ . (2 ) at thls temperatore.

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NEWTON AND SWADDLE: COBALT(II1) AMMINES 2757

C O ( N H , ) ~ ~ + and CO(NH,),OH,~+ in dilute the same relative amounts whether the starting perchloric acid are Co(H20),2f, N H 4 + , N2, material be the hexaammine or the penta- N,O, and a relatively small amount of O,, and ammine, which supports the conclusion that the that the yield of N,O relative to N, is roughly initial step in the deconlposition of the former the same for both complex ions and close to that complex is aquation to the latter without redox. expected if NH,OH were the intermediate giving A minor redox pathway may conceivably com- rise to both these gases. Pete with aquation in the case of the tetraammine

(2) but even so the important redox step in the Discussion

The first step in the thermal decomposition of CO(NH,),~+ in acidic aqueous solution is loss of ammonia to yield CO(NH,),OH,~ +. The sub- sequent aquation of this aquo-species and of its successor CO(NH,),(OH,),~+ are much more rapid than that of the hexaammine because paths with inverse dependence on the hydrogen- ion concentration become accessible when a coordinated aquo group is present (eq. 2), and these conjugate base pathways are favored over the acid-independent aquation paths at the acidi- ties used in these experiments, as the data of Table 4 will show.

Interestingly, the rate coefficient for the acid independent path (k,) is essentially the same for CO(NH,),~+ and Co(NH,),OHZ3+, and if any- thing somewhat less for c ~ s - C O ( N H ~ ) , ( O H , ) , ~ ~ ; this is precisely the pattern found for the stepwise aquations of the corresponding chromium(ll1) ammines (17). However, whereas the overall decomposition rates of the cobalt(II1) complexes are greatly increased by the incursion of con- jugate base pathways which increase in impor- tance as the number of aquo groups increases, conjugate base pathways make no significant contribution to the aquation rates of any of the Cr(II1) ammines from C ~ ( N H , ) , O H , ~ + to Cr(NH3)(OH,),3 ' in 0.1 M HCIO, (1 7). This is another illustration of the great susceptibility of Co(II1) systems to conjugate base hydrolysis mechanisms, relative to their Cr(II1) analogs, a phenomenon well documented for hydrolyses in alkaline media (18). The result is that the Cr(1II) ammines aquate via a series of ever-slower steps, resulting in complicated rate laws (1 7), whereas for the hexa-, penta-, and tetra-amminecobalt- (111) ions the overall rates of successive steps become faster and faster, so that spectral changes are very nearly first-order kinetically.

On reaching C O ( N H , ) , ( ~ H , ) ~ ~ +, the decom- position of Co(II1) ammines proceeds by direct redox to Co(lI), N,, N,O, and a minor amount of 0 , . The gaseous products appear in essentially

sequence is clearly the decomposition of Co- (NH3)3(OH2)33+ (3).

The redox decomposition of Co(NH,),- + to cobalt(I1) must involve a single-elec-

troll transfer a t some stage, and oxidation of a water molecule (or aquo ligand) to O H or of an ammine ligand to NH, afford the only dlrect one-electron initial redox steps. Since both N,O and a minor amount of 0, are ultimately formed, in addition to N,, an oxygen-containing radical must be involved in the overall redox process. This radical is almost certainly OH, and, because NH, is not expected to produce O H by abstrac- tion of H fro111 H 2 0 (the H-OH bond strength being 1 19 kcal m o l l as against 103 for H-NH, (19)). we may infer that O H is the,first radical to be formed, as in eq. 4.

It is known (14) that O H does not attack coordinated NH, ligands. Thus, as long as the assemblage of species on the right hand side of eq. 4 remains intact within the solvent cage, further reaction (as distinct from recombination) can only occur following dissociation of one of the ammine ligands from Co(NH3)3(OH2)32t.

The pseudo first-order rate coefficient for reaction 5 is about 6 x lo4 s- ' at 25" (20), and will be of the order of 10' s - ' at the tempera- tures prevailing in the present study. Reaction 5 will therefore be slower by at least three orders of magnitude than the diffusion of H + out of the solvent cage, a process which would inevitably be fast by operation of the Grotthus chain mechanism and which in this case will be accel- erated and rendered essentially irreversible by the double positive charge on the complex ion. Con- sequently, as long as O H remains in the solva- tion cage, reaction 6 (for which the bimolecular

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2758 C A N . J . CHEM. VOL. 5 2 . 1974

rate coefficient is 1.0 x lo8 M-I s-I at 25" (21)) can occur without interference from protonation of the freshly-released NH,, a significant point, since OH does not attack NH4+ (21). Alterna- tively, NH, might be produced by direct oxida- tion of an NH, ligand in CO(NH,),(OH,),~+ by Co(lII), although, as noted above, coordination deactivates NH, toward oxidation (14).

Thus, the free-radical product which emerges from the solvent cage will be OH itself or NH,, which, being odd-electron species, will survive in bulk solution until meeting another free radical. If the emergent species is NH,, it will scavenge an OH radical which either has been freshly formed as a Co(I1)-OH pair or has escaped from the solvent cage of its parent complex ion. In either case, the product will be hydroxylamine (eq. 7); an encounter between NH, and another NH, is less likely, as OH may well be the pre- cursor of NH, (eq. 6), and reaction 7 is known to be rapid (bimolecular rate coefficient 9.5 x lo9 M-' s- ' at 25" (21)).

It is well known that hydroxylamine decom- poses thermally in acidic aqueous solution to give N,, N 2 0 , and NH4+ (22), and we have established that the relative amounts of N, and N,O observed for the thermal decompositions of N H 2 0 H , Co(NH,),,+, and CO(NH,),OH,~ +

are sufficiently similar to indicate that hydroxyl- amine (or rather the hydroxylammonium ion) is indeed the first major non-radical oxidat~on product in the decomposition of the Co(II1) ammines to Co(I1).

Alternatively, an OH radical may occasionally escape from the solvent cage before reaction 5 occurs and scavenge another O H radical, similarly formed, rather than NH,. In this case, hydrogen peroxide will be formed instead of N H 2 0 H , and will decompose to O,, which is indeed observed as a minor reaction product.

The stoichiometry of the overall thermal de- composition of CO(NH,),~' may be summarized by eqs. 10-12, in which [ l l ] is most important and [12] least.

The kinetics and products of the thermal de- compositions of Co(NH,),,+ and Co(NH,),- OH,,' in acidic aqueous solution bear a marked resemblance to those of the corresponding decompositions in molten NH4HS04 and in 9 7 z sulfuric acid (1 1, 23). However, in the latter solvents, the intermediates are inevitably bisul- fato rather than aquo complexes, and the gaseous product is exclusively N, (apart from trace amounts of NO and SO,), as might be expected, since there is no obvious source of OH radicals to give rise to N,O or 0,. The enthalpies of acti- vation associated with the loss of the first NH, ligands from the hexaammine, pentaammine, and cis-tetraammine in 97% H2S04 are in the range 38-42 kcal mol-I, that is, very close to those for the corresponding processes in water; this suggests that the activation process is inde- pendent of the nature of the solvent, i .e . , that it is dissociative. In contrast, the initial step in the thermal decomposition of solid salts of Co(II1) ammines is evidently redox to Co(II), rather than ligand substitution (24-26).

In the highly acidic solvent H2S04, conjugate base pathways are unlikely to be of any impor- tance, so that the rates of loss of NH, from CO(NH,),HSO,~+ and Co(NH,),,+ are gov- erned by k , alone and are therefore closely similar. Thus, Sutula and Hunt (1 1 ) observed substantial accumulations of the pentaammine complex as the hexaammine underwent solvol- ysis in 97'7, H2S04, and the kinetics of decom- positions in H 2 S 0 4 were more complicated than those in water, in which the conjugate base path- ways facilitate removal of reaction intermediates.

We thank Dr. A. W. Boyd for discussions, and the National Research Council of Canada for financial support.

1. W. E. J O ~ E S , L. R. CAREY, and T. W. SWADDLE. Can. J. Chem. 50, 2739 (1972).

2. W. D. STA\LEY, S. LUM, and C. S. GARNER. J. Inorg. Nucl. Chem. 35, 3587 (1973).

3. S. LUM, W. D. STANLEY, and C. S. GARNER. J. Inorg. Nucl. Chem. 35, 1293 (1973).

4. W. D. STANLEY, T. DAVIES, T. D. TULLIUS, and C. S. GARNER. J. Inorg. Nucl. Chem. 35, 3857 (1973).

5. J. D. WHITE, J. C. SULLIVA~, and H. TAUBE. J. Am. Chem. Soc. 92, 4733 (1970).

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6. J. BJERRUM and J. P. MCREYNOLDS. Inorg. Synth. 2, 18. F. BASOLO and R. G. PEARSON. Mechanisms of 218 (1946). inorganic reactions. 2nd Ed. J. Wiley and Sons, Inc.,

7. W. E. JONES and T. W. SWADDLE. Can. J. Chem. 45, New York, N.Y. 1967. p. 188. 2647 (1967). 19. Handbook of chemistry and physics. 51st Ed. The

8. R. J. MUREINIK, A. M. FELTHAM, and M. SPIRO. J. Chemical Rubber Company, Cleveland, Ohio. 1970. Chem. Soc. Dalton, 1981 (1972). p. F164.

9. R. E. KJTSEN. Anal. Chem. 22, 664 (1950). 20. M. SIMIC and J. LILIE. J. Am. Chem. Soc. 96, 291 10. G . DAVIES and K. 0. WATKINS. J. Phys. Chem. 74, (1974).

3388 (1970). 21. P. B. PABSBERG. R i s ~ Report No. 256. Danish 11. R. A. SUTULA and J. B. HUNT. Inorg. Chem. 11, Atomic Energy Commission, Rim, Roskilde, Den-

1879 (1972). mark. 1972. 12. H. C. HELGESON. J. Phys. Chem. 71, 3121 (1967). 22. K . JONES. In Comprehensive inorganic chemistry. 13. D. KATAKIS and A. 0. ALLEN. J. Phys. Chem. 68, Vol. 2. Edifrd by A. F. Trotman-Dickenson. Per-

1359 (1964). gamon Press, London. 1973. p. 272. 14. D. THUSIUS and H. TAUBE. J. Phys. Chem. 71, 3845 23. J. A. DUFFY and W. J. MACDONALD. J. Chem. Soc.

(1967); J. Am. Chem. Soc. 88, 850 (1966). (A), 1160 (1971). 15. D. A. BUCKINGHAM, A. M. SARGESON, and L. G. 24. W. W. WENDLANDT and J. P. SMITH. The thermal

MARZILLI. J. Am. Chem. Soc. 89, 3428 (1967). properties of transition metal ammine con~plexes. 16. L. G. S I L L ~ N and A. E. MARTELL. Stability constants. Elsevier Publishing Co., Amsterdam. 1967.

Special Publication No. 17. The Chemical Society, 25. A. V. ABLOV and T. A. MAL'KOVA. RUSS. J. Inorg. London. 1964. p. 55. Chem. 14, 1594 (1969).

17. G. GUASTALLA and T. W. SWADDLE. Inorg. Chem. 26. E. L. SIM~IONS and W. W. WENDLANDT. J. Inorg. 13, 61 (1974). Nucl. Chem. 28, 2187 (1966).

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