Oxidation of hydrocarbons. 12. Kinetics and ...

Oxidation of hydrocarbons. 12. Kinetics and mechanism of the oxidation of trans-cinnamic acid by ruthenate and perruthenate ions' DONALDG . LEE' AND STUARTHELLIWELL Department of Chemistry, University of Regina, Regina, Snsk., Canada S4S 0A2

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Received May 30, 1 9 8 3 ~ DONALD G . LEEand STUART HELLIWELL. Can. J . Chern. 62, 1085 (1984). Recent theoretical calculations by RappC and Goddard have pointed to the importance of spectator 0x0 groups in stabilizing intermediates produced during thc reduction of high-valent transition metal oxides; intermediates which contain triply-bonded metal 0x0 groups are predicted to be unusually stable. In an attempt to test these conclusions experimentally, the rates of oxidation of cinnamate ion by ruthenate and permthenate ions have been studied. Since the most likely intermediate formed from permthenate ion (but not ruthenate ion) would contain a stabilizing, triply-bonded spectator 0x0 group, it would be expected to react faster if the calculations are correct. In agreement with these predictions, it was found that the activation energy (AG*) for the ruthenate reaction was about 7 kcal/mol higher than for the corresponding permthenate reaction. DONALD G. LEEet STUARTHELLIWELL. Can. J . Chem. 62, 1085 (1984) Les recents calculs thCoriques de RappC et Goddard ont mis en Cvidence I'importance des groupes 0x0 spectateurs dans la stabilisation dcs intermediaires obtenus lors de la riduction d'oxydes de mCtaux de transition de valence ClevCe; d'apres ces calculs, les intermCdiaires qui contiennent les groupes 0x0 triplement liCs au metal seront particulierement stables. Dans le but de verifier expCrimentalement cette hypothkse, on a CtudiC les vitesses d'oxydation de I'ion cinnamate par les ions ruthenate et permthCnate. Puisque l'intermtdiaire le plus communtment form6 a partir de I'ion permthenate (mais non B partir de I'ion ruthhate) contiendrait le groupe spectateur 0x0 triplement liC, on pourrait s'attendre 21 ce qu'il soit plus rkactif si les calculs sont corrects. En accord avec ces prkvisions, on a trouvC que I'tnergie d'activation (AG*) de la rCaction du ruthhate est d'environ 7 kcal/mol plus ClevCe que celle de la rCaction correspondante du permthCnate. [Traduit par le journal]

Introduction Transition metal oxides and oxyanions have been used extensively as oxidants in both inorganic and organic chemistry (1). The driving force in most such reactions is generally considered to be a lowering of the oxidation number of the oxidant, which may formally be considered as a flow of electrons into the d orbitals of the metal. Since many of the high-valent do, d l , and d2 oxidants are capable of accepting several electrons before reaching a stable oxidation state, these reactions usually proceed in several steps with the formation of a number of intermediates of greater or lesser stability. The reactions are consequently of interest from both a practical and a theoretical point of view. RappC and Goddard (2) have recently described results, obtained from a theoretical study, that are extremely useful for predicting which intermediates are most likely to be formed during the course of several of these reactions. By use of ab initio methods they found that the metal-0x0 bonds present substantially influenced the stability of an intermediate. Metal-0x0 groups which have triple bond character were found to impart greater stability to the intermediate than those metal-0x0 groups which were doubly bonded. The metal-0x0 double bonds are formed by the overlap of ad,~ orbital on the metal with a p, orbital on oxygen to give a o-bond (each element contributing one electron) and the parallel overlap of a dxzorbital on the metal with a p , orbital on the oxygen to give a T-bond. The metal-0x0 triple bonds are formed, however, by the overlap of 4zand d,,?orbitals on the metal withp, andp? orbitals on the oxygen to give two T-bonds along with a third bond formed by the overlap of a filled p, orbital on the oxygen with an empty d22on the metal. Since the triple bonds lend added stability to the structures under consid-

' For part 1 1 see ref. 29. 'Author to whom correspondence may be addressed. 'Revision received January 10, 1984.

eration, they would be formed whenever the appropriate orbitals were available. In order to identify these bonds in structural formulas, the convention of using two solid and one broken line was adopted; i.e., M=O (2). As an example of the importance of the nature of these spectator 0x0 groups, RappC and Goddard pointed to the difference in the energy changes which take place as reactions [l] and [2] occurred. Reaction [l] is endothermic because it involves destruction of a metal-0x0 triple bond, whereas reaction [2] is exothermic because it involves formation of a metal-0x0 triple bond.

Spectator 0x0 groups with triple bonds can form only when two d orbitals are available for bonding to a single oxygen. In the case of chromyl chloride, 2, a triple bond is not possible because use of both d orbitals for one bond would deprive the second oxygen of its T-bond. The bonding in chromyl chloride is thus analogous to the bonding in carbon dioxide, while the bonding in 1 is similar to that in carbon monoxide. As a further example of the importance of this concept, RappC and Goddard pointed to the difference in the reactions of chromyl chloride and permanganate ion with alkenes (2). Both oxides react with carbon-carbon double bonds to give the corresponding metallocyclooxetanes, 3 and 4 (3). However, 4 goes on to form the corresponding cyclic manganese(V) diester, 5 (and subsequently to yield diols (4) or cleavage prod-

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LEE A N D HELLIWELL

1087

TABLE 3. Isotope effects on the initial reaction between cinnarnate and ruthenate ions (reaction [9], Schernc I )

Substrate"

k z ( M - ' s-' x 10')"

kt~lkr,

Cinnarnate ion Cinnarnate-a-d Cinnarnate-P-d

5.94 2 0.15 7.33 + 0.09 7.02 + 0.20

0.81 + 0.03 0.85 r 0.04

"[Substrate] = 1.21 X lo-' M ; [NaOH]

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'Temperature

Frc. 1 . Dependence of k2 on base concentration; T = 85°C.

the actual oxidant (9). Occurrence of such a reaction would lead to a kinetic rate law that was initially second order in [RUO;-1. Work from several laboratories has indicated that reactions between high-valent transition metal oxides and alkenes proceed via organometallic intermediates which involve the formation of a T-bond between the carbon-carbon double bond and the metal (2, 3, 10). These T-complexes then rearrange to metallocyclooxetanes and eventually to cyclic diesters as in eq. [8]. In the case of ruthenate it is expected that formation of the metallocyclooxetane, 8, would be quite rapid since it involves

X RuOi-

/

k t

\

66:--1.---

&-(jy- - - --

J=Q~-

I 06-

fast

H B

& 0-RU

/ \o-

-0

8

formation of a ruthenium(V1) compound with a triply-bonded spectator 0x0 (2). The slow step of the reaction sequence depicted in eq. [8] would therefore be the conversion of 8 into the cyclic ruthenium(1V) diester, 7. Since this step results in a conversion of ruthenium from a d 2 to a d 4 state, it would be expected to be slower than the previous steps which involve no change in oxidation state. The inverse secondary isotope effects that are observed when the olefinic hydrogens are sequentially replaced by deuterium (Table 3) establish that the a and P carbons have rehybridized from sp2-p in the ground state to sp3 in the transition state (1 1) and are consistent with the occurrence of the transition state

=

=

0.532 M .

85.0°C.

between structures 7 and 8. The unfavorable entropy of activation (AS = -30.6 + 1.3 eu) is also consistent with a highly structured transition state (12). The activation entropy for the reaction (Table 4) is, in fact, very similar to that for the oxidation of cinnamate ion by permanganate, which is known to proceed by way of a cyclic manganate(V) diester (13). The larger enthalpy of activation for the ruthenate reaction is probably a consequence of the fact that the cyclic diester, 5, formed in the permanganate reaction (eq. [4]) is stabilized by the presence of a spectator 0x0 while the ruthenium(1V) diester, 7, is not. The first-order dependence of the rate of reaction on the concentration of ruthenate ion does not extend throughout the entire course of the reaction, as shown by the curvature in the plot of In [RUO:-] vs. time (Fig. 2). This deviation from linearity is indicative of secondary reactions of the initially formed products. Although this could potentially lead to an unwieldly large number of possible reaction schemes, all but one reasonable possibility could be eliminated by the use of computer modeling. Employing a previously published program (14), we examined 28 possible reaction sequences by computing the rate constant for each step that is necessary to give the best fit between experimental and theoretical absorption values at various time intervals. Although several reaction schemes gave excellent fits for a particular set of conditions, only one sequence of reactions gave rate constants that were the same regardless of the initial reactant concentrations. This reaction sequence has been summarized in Scheme 1. The ability of this sequence to reproduce the experimental absorbance values is indicated by the data in Table 5. The first step of this scheme is assumed to be a reaction between the unsaturated carboxylic acid and ruthenate to accommodate the previously described observation that the initiaI rate of reaction is first order in both Ru0:- and cinnamate ion. Since the first step was also observed to be inversely dependent on base concentration, it has been shown in Scheme 1 as being subject to acid catalysis. The value for k? could be determined directly from initial rate studies (Table 2); however, the magnitude of k3 could not be exactly specified since any value large enough to maintain a steady-state concentration of the intermediate 7 was adequate. For the purpose of the computer simulation a rate constant of 80 M-' s-' was used. The values k5 and k6 were also assumed to be much larger than k2 or k, and were arbitrarily set at 50 M-' s-'. Reactions [12] and [13] are the product-forming steps and in all likelihood refer to the oxidation of benzaldehyde and glyoxylate ion (which would be formed by the oxidative cleavage of the double bond) to give the corresponding carboxylic acids. Since such reactions were qualitatively found to be extremely fast under these conditions, the assumption that k5 and k6 would be large seems justified. The structure of intermediate 9 is open to more speculation; however, at least one attractive possibility that is consistent

CAN. J . CHEM. VOL. 62. 1984

TABLE 4. Activation parameters for the reaction of cinnamate ion with MnO,, RuO,, and RUO:Oxidant MnO, MnO, RuO, RUO;-

AH' (kcal/mol)

Conditions

0.99 M 0.50 M 0.39 M 0.54 M

HCIO, NaOH NaOH NaOH

4.2 + 4.7 2 5.0 2 12.5 ?

0.1 0.1 0.1 0.5

AG (kcal/mol)"

ASf (eu) -30 -29 -32 -31

2

+ 2 2

1 1 1 2

13.1 13.3 14.5 21.7

?

2 2 ?

Temp. range

0.4 0.4 0.4 1.1

12-41°C 0-40°C 7-45°C 65-95°C

"Actcalculated at 25'C.

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b ~ l plots l of In k / T vs. 1 / T exhibited correlation coefficients of 0.998 or better for a minimum of 5 temperatures.

TABLE5. Observed and calculated absorbances for the oxidation of cinnamate ion by sodium ruthenate Time (5 x 10')

Absorbance (obs.)

Absorbance (calcd.)

0 1.83 5.09 8.17 12.4 16.8 21.2 25.8 32.0 39.4 49.8 61.2 75.4 93.0

0.803 0.778 0.726 0.686 0.633 0.587 0.545 0.507 0.464 0.417 0.365 0.322 0.282 0.244

0.803 0.775 0.728 0.687 0.637 0.590 0.549 0.510 0.464 0.417 0.364 0.320 0.282 0.256

R

[9]

#:

+ ~~0425 HWX slow

H

R

[lo]

H

H++X

R

o\Ru,o

R

+ RUO:-

[''I

o\ Ru-0P

\o-

R

H

'%q?(lo Gx

H

HHX + RUO:I

fast

o\Ruo, HO/

H

0 " 0

II

k4

dowest t RCH

O 'H

0

1I

+ XCH + HRuO74

Ru-0-

o& il

[12]

RCH

[13]

XCH

I1

ks + ~~042+ RCO? + HRu034

+ ~~042- k6

t XCOT

+ HRuOj- 4

X = COT

SCHEME1

TIME (min) FIG.2. Attempted first-order plot for the oxidation of cinnamate by ruthenate. [RUO:-] = 4.37 X M ;[cinnamate] = 5.71 x lo-' M ; [NaOH] = 2.46 M ;T = 85°C. with all of the experimental evidence is available if 9 is assumed to be a cyclic ruthenate(V1) diester formed by the oxidation of 7. Such a two-electron transfer reaction is consistent with the known properties of ruthenate under these conditions (15). Furthermore, it would not be unreasonable to observe a large rate constant for such a reaction; many similar electron

transfer processes between transition metal oxyanions are known to occur rapidly (16). Development of a triply-bonded spectator 0x0 group as the reaction progresses would also serve to increase its rate (2). Intermediate 9 could undergo one of two possible reactions: (i) it could hydrolyze to give a diol, or (ii) it could undergo oxidative decomposition to give cleavage products. Since only compounds resulting from the cleavage of the carbon-carbon double bond could be found among the products of the reaction, it may be assumed that 9 decomposes as depicted in eq. [Ill. - other evidence consistent with this suggestion includes the normal secondary kinetic isotope effects (Table 6) which are indicative of a rehybridization of the carbons from sp3to sp2 as the reaction proceeds (1 l), and a favorable entropy of activation (AS' = 7.7 ? 2.1 eu) as would be expected for a fragmentation reaction (12). A large enthalpy of activation (31.0 + 0.7 kcal/mol) for this step is indicative of similar reactions in which a-bonds are cleaved (12, 13). Finally, when the concentrations of the reactants and intermediates are calculated from the computer model for Scheme 1 and plotted against time it is

LEE AND HELLIWELL

TABLE^. Isotope effects on reaction [ I l l of Scheme I

Substrate"

k2 (s-' x 10')

Cinnamate ion Cinnamate-a-d Cinnamate-P-d

1.94 ? 0.01 1.37 0.04 1.65 ? 0.07

*

k~lkn 1.42 1.18

* 0.05 ?

0.12

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"[Substrate] = 1.21 X 10.' M ;[NaOH] = 0.532 M "Temperature = 85.O"C.

WAVELENGTH nm

FIG.4. Sequential scans as the oxidation of cinnamate by ruthenate proceeds. Spectra taken at 0, 196, 634, 1320, 2320, and 5380 s (top to bottom).

2000

6000

TIME

(5)

FIG. 3. Calculated concentrations of reactants, intermediates, and products for a reaction proceeding according to Scheme 1.

observed that 9 is the only intermediate which accumulates as the reaction proceeds (Fig. 3). Since this intermediate is in the same oxidation state as RUO;-, it is not to be expected that its presence would produce detectable changes in the uv-vis spectrum as the reaction progresses (Fig. 4); oxyanions of a metal with the same oxidation state often have very similar spectra (17). 'The rate law for the process summarized by the reactions in Scheme 1 may be derived by noting that reaction [lo] does not result in a net change in oxidation state, the property which presumably controls the spectral changes that were used to monitor the rate of reaction. Therefore the experimental rate law would be given by eq. [14].

In the initial stages of the reaction, [9] will be vanishingly small; i.e., when [RUO;-] = [RUO;-]~,~,~,~, [9] = 0. Consequently, when initial reaction rates are considered, the last term in eq. [14] can be neglected and the rate law for the initial reaction is therefore first order in ruthenate, first order in cinnamate, and inversely dependent on base concentration. This is in accordance with the experimental observations (see Fig. 1 and Tables 1 and 2). As [9] increases, this approximation cannot be made and the rate expression becomes one obtained for consecutive reactions of mixed order (18). Such expressions are not amenable to treatment by classical methods and are best handled by numerical methods as has been done in this instance (18). It may be noted, however, that the second term in eq. [14] will become dominant as the concentration of 9 increases (Fig. 3). When this happens one would expect the reaction to become second order in ruthenate since two moles of Ru0;- are consumed in the formation of 9. As a consequence, deceptively

TIME ( m i n !

FIG. 5. Pseudo second-order plot for the oxidation of cinnamate by ruthenate. [RUO:-] = 4.37 X lo-' M; [cinnamate] = 5.71 x lo-' M; [NaOH] = 2.46 M; T = 85°C.

good pseudo second-order plots are obtained when the oxidation of an excess of cinnamate is studied under non-initial rate conditions (Fig. 5). Perruthenate oxidations Because the reaction of perruthenate ion with cinnamate ion under alkaline conditions is so much faster than the corresponding reaction with ruthenate ions, it was not possible to study the two reactions at the same temperature or in exactly the same way. Whereas it was convenient to follow the rate of ruthenate oxidation at 85OC by use of conventional spectrophotometric methods, the rate ofthe permthenate oxidations could only be followed at a lower temperature (21°C) and with the aid of a stopped-flow system. The reaction was found to be first order in perruthenate (Fig. 6), first order in cinnamate ion (Fig. 7), and independent of base concentration (Table 7). As the reaction proceeded, the green permthenate solution was rapidly changed into a bright orange solution that exhibited the same uv-vis absorption

CAN. 1. CHEM. VOL. 62. 1984

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TABLE 7 . Effect of base concentration on the rate of oxidation of cinnamate ion by sodium perm thenate"

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I

I 10

I

I 10

I

I

k ( M - ' s-I)

0.040 0.080 0.140 0.200 0.280 0.350

190 199 196 191 179 177

"[Substrate] = 3.50 x temperature = 21.0DC;(L

I

30

lo-' =

M;

0.500

M.

TIME Isl

'FIG. 6 . Pseudo first-order plot for the oxidation of cinnamate by M ; [cinnamate] = 3.50 x permthenate. [RuOh] = 2.10 x M; [NaOH] = 0.350 M ; T = 21°C.

[NaOHl (MI

TABLE8. Isotope effects on the rate of oxidation of cinnamate ion by sodium perruthenate" Substrate

[NaOH] ( M )

Cinnamate ion (1.75 X lo-' M ) Cinnamate-a-d ( I .75 x lo-' M ) Cinnamate ion (3.40 X lo-' M ) Cinnamate-P-d (3.40 x lo-' M ) "Temperature

=

21 .O°C;(L

=

0.098 0.098 0.282 0.282

k ( M - ' s-I)

215 t 222 133 & 138?

6

*6 2 3

0.500 M.

The lack of a substantial secondary deuterium isotope effect for this reaction sequence (Table 8) indicates that the carbon atoms are sp2-p hybridized in the transition state. Consequently it appears as if the rate-determining step is likely oxidative decomposition of the cyclic ruthenium(V) diester as in eq. [16]. In this respect the mechanism is parallel to the one proposed above for RUO;- oxidations, where decomposition of a similar cyclic diester was also found to be the slowest step in the reaction sequence.

FIG. 7 . Dependence of the rate of permthenate oxidations on cinnamate concentration.

spectrum as ruthenate solutions. These solutions reacted more slowly to give ruthenium dioxide and benzoic acid. The activation parameters for this reaction were found to be very similar to those observed for the corresponding reaction of permanganate with cinnamic acid (Table 4). Hence it appears as if the reaction likely involves an initial addition of the oxidant to the double bond, giving a ruthenium(V) cyclic diester, 10, as in eq. [15]. However, since there was no indication of the formation of inorganic products other than RUO$- the intermediate, 10, must undergo a subsequent reaction leading to the formation of ruthenate. A likely possibility is that 10 could, in analogy with the corresponding cyclic manganese(V) diester, oxidatively decompose to give cleavage products and ruthenate(II1) as in eq. [16]. The ruthenate(II1) could then react with perruthenate to give ruthenate(V1) as in eq. [17].

H

0

R'

-+

[I61 *R

II

RCH

0

II + RCH + RUOF

0,9!h Ru

o& \o-

The difference in activation energies (AAG') for the oxidation of cinnamate by ruthenate and perruthenate is about 7 kcal/mol. From this it can be calculated that the rate of the

LEE A N D HELLIWELL

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perruthenate reaction is about 5 X lo5 times a s fast a s the corresponding ruthenate reaction at room temperature. Although part of this effect may b e d u e to the higher reduction potential of perruthenate, a large portion must b e attributed to a difference in the relative stabilities of the transition and ground states as described above. T h e fact that ruthenate, although a vigorous oxidant for alcohols, reacts very slowly with carbon- carbon double bonds has made it an excellent reagent for the oxidation of unsaturated alcohols (19). Several examples of its use for this purpose (20-22), a s well a s for the oxidation of certain saturated alcohols (23), have appeared in the recent literature.

Experimental Sorliutn ruthenate (15. 24), ~, Ruthenium tetroxide (25) was first prcpared by reaction of 5.0 g of hydrated ruthenium dioxide with 350 mL of commercial bleach (5.25% sodium hypochlorite). After stirring for approximately 10 min, the RuO, which had formed was extracted into carbon tetrachloride (5 X 75 mL) which in turn was washed with 3 X 100 mL distilled water to remove excess hypochlorite. An equal volume of I .O M NaOH was then added to the carbon tetrachloride solution and the heterogeneous mixture stirred for several hours (until the organic layer was colorless) and the deep orange basic ruthenate solution separated. Bcfore being used in a kinetic experiment these solutions were heated at 75OC for several days. This ensures that all of the ruthenium is present in the + 6 oxidation state and that the last traces of sodium hypochloritc have been removed. The purity and concentration of these solutions were thcn assayed using procedures described in the literature (24). Sodium perruthenate (24, 26) This reagent was also prepared from the reduction of ruthenium tetroxide. After the oxidation of hydrated ruthenium dioxide by bleach, as described above, a stream of air was used to carry the RuO, over into 0.3 M NaOH cooled at -3OC. The sodium permthenate thus formed was allowed to stand at O°C overnight to ensure complete reduction of ruthenium tetroxide to sodium permthenate. Because of the instability of permthenate solutions, these were madc up as required. Product studies A solution of 4.5 X lo-' M RUO;- (170 mL of 1.0 M NaOH) was added to 0.30 g of cinnamic acid dissolved in 50 mL of 0.5 M NaOH. The reaction was allowed to proceed overnight at 6OoC. The black ruthenium dioxide was filtered off and washed with warm distilled water. The basic solution was extracted with 3 X 75 mL portions of ether, saturated with sodium chloride, and acidified with dilute hydrochloric acid. The resulting solutions were extracted with 3 X 75 mL portions of ether. Evaporation of the ether after drying with anhydrous MgSO, left a white residue from the acid extract but nothing from the basic extract. Identification of the white residue indicated that it consisted of benzoic and some unreacted cinnamic acid. The yields based on unrecovered starting material were 99.0 and 99.5% benzoic acid for duplicate runs. he-oxidation of fumaric acid was carried out in a similar manner up to and including the filtering to remove ruthenium dioxide. After neutralizing the filtrate with hydrochloric acid, 10 mL of 10% CaC1, solution was added. This produced an immediate white precipitate of calcium oxalate. The solution was allowed to stand overnight at 70°C to ensure complete precipitation and coagulation of the calcium oxalate (27). The precipitate was then filtered through a sintered glass crucible, dried at 105OC. and weighed. Powder X-ray photographs taken of both this precipitate and a prepared sample of calcium oxalate were shown to be the same. Samples of the precipitate were then analyzed by dissolving a sample (about 0.2 g) in 20 mL of 1 : 8 H,SO,, diluting it with 70 mL of distilled water, heating to 80-90°C, and titrating with standard KMnO, solution (28). Yields of 88 and 95%

1091

oxalic acid were obtained for duplicate experiments. When perruthenate ion reacted with cinnamic acid at O°C, the deep yellow-green RuO, color was instantly replaced by an intense orange color typical of ruthenate ion. The ruthenate color then was observed to disappear as ruthenium dioxide formation took place. Hence it was decided to attempt to quench the ruthenate formed with sodium arsenite in the following way. To 100 mL of 0.3 M NaOH (at O°C) containing 0.40 g of cinnamic acid was added 175 mL of sodium permthenate solution (0.025 M). Immediately following this addition, 125 mL of 0.1 M NaAsO' were added. However, due to the rapidity of the reactions, quenching was only partially successful and only a qualitative estimation of the products of the initial reaction could be made. Both benzaldehyde and benzoic acid were isolated after a work-up as described above. A more quantitative result was obtained when the reaction was allowed to go to completion without quenching. A 99% yield of benzoic acid was then obtained. -

~

Kinetic studies In the kinetic studies of the oxidation of cinnamic acid by both ruthenate and permthenate, pseudo order techniques were used throughout; i.e. greater than a ten-times equivalent excess of substrate over oxidant was used. It was found unnecessary to maintain a constant ionic strength for the ruthenate oxidations because varying the ionic strength had little or no effect on the rate of reaction. However, since the rates of the perruthenate reactions were not independent of the ionic strength, it was maintained at a p = 0.500 M by use of sodium perchlorate for these reactions. In the case of ruthenate oxidations, solutions of the desired base concentrations were made up (2-L batches) and stabilized at 75OC for 2 days. The reactions were then carried out in 125-mL ground glassstoppered erlenmeyer flasks thermostated in an oil bath at 85OC. In a typical set of experiments, eight flasks, each containing 100 mL of ruthenate solution. were allowed to thermostat for 1 h before the desired amount of substrate was added to start the reaction. Aliquots were removed at appropriate intervals and then quenched by cooling in ice. Absorbance rcadings were made using a Perkin-Elmer 356 dual wavelength, double beam spectrophotometer. This instrument has the advantage that a reference beam compensates for any precipitate which may have formed during the reaction. The wavelengths used were 460 and 570 nm. Little absorption by ruthenate occurs at 570 nm, hence this wavelength was used as the reference beam. The absorbance readings were converted into ruthenate concentrations using a calibration curve based on ruthenate solutions of known concentrations. In several experiments the aliquots were first cooled, then centrifuged to remove ruthenium dioxide and scanned from 300 to 600 nm in an attempt to detect possible intermediates. As the plots in Fig. 4 indicate, there was no evidence for the formation of long-lived intermediates. Time vs. absorbance plots obtained in this way were very similar to those obtained using the dual wavelength mode. The reaction between permthenate ion and cinnamate ion could be followed spectrophotornetrically by monitoring the disappearance of permthenate ion at 385 nm. Because of the rapid nature of the reaction, all experiments were performed using a Durmm Model D-110 stopped-flow spectrophotometer. In a typical experiment, 0.5 mL of RuO, solution (ca. 0.05 M ) was pipetted into a 125-mL erlenmeyer flask containing approximately 50 mL of a NaOH solution and a magnetic stirring bar. The flask was stoppered and the solution degassed by reducing the pressure to approximately 95 Torr and stirring. When all of the RuO, had been reduced to RuOh (determined spectrophotometrically). one of the stopped-flow syringes was filled with this solution and the other with a solution of the substrate. Equal volumes of the two reactants were forced through the mixing jet and into the cuvette where the absorbance measurements were made. The resulting kinetic curve, traced on an oscilloscope, was then photographed. Determinations were usually carried out in triplicate. In order to check for the formation of possible intermediates, kinetic curves were run at a number of different wavelengths and spectra synthesized as a function of time (Fig. 8). As may be seen from the

1092

CAN. I . CHEM. 1

9. 10. 11.

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12.

I

I 400

I 500

I

450

WAVELENGTH

I

650

13.

nm

FIG. 8. Sequential scans as the oxidation of cinnamate by perruthenate proceeds. Spectra taken at 0.5-s intervals, top to bottom. good isosbestic point, no intermediate could be detected, the spectra after 5 s having a close resemblance to that of ruthenate.

14. 15.

Acknowledgements

16.

Financial assistance from the Natural Sciences and Engineering Research Council of Canada, the donors of the Petroleum Research Fund, administered by the American Chemical Society, and the Committee o n Higher Education, Imperial Oil Limited is gratefully acknowledged. 1. R. STEWART.Oxidation mechanisms. Benjamin, New York. 1964; K. B. WIBERG.Oxidation in organic chemistry, Part A. Academic Press, New York. 1965; W. S. TRAHANOVSKY. Oxidation in organic chemistry, Parts B, C, and D. Academic Prcss, New York. 1973, 1978, and 1982; D. ARNDT.Manganese compounds as oxidizing agents in organic chemistry. Open Court, La Salle. 1981; D. G. LEE.The oxidation of organic compounds by permanganate ion and hexavalent chromium. Open Court, La Salle. 1981. 111. J. Am. Chem. Soc. 104, 2. A. K. RAPPEand W. A. GODDARD 448 (1982); 104, 3287 (1982); A. K. RAPPE.Ph.D. Thesis, California Institute of Technology. 1980. J. A. Y. TERANISHI, and J. E. BACKVALL. 3. K. B. SHARPLESS, Am. Chem. Soc. 99, 3120 (1977). 4. D. ARNDT.Manganese compounds as oxidizing agents in organic chemistry. Open Court, La Salle. 1981. pp. 207-212. 5. R. N. HASZELDINE. J. Chem. Soc. 4291 (1955); 4302 (1955); and J. E. OSBORNE.J. Chem. Soc. 3880 R. N. HASZELDINE (1955); R. N. HASZELDINE and K. LEEDHAM. J. Chern. Soc. 1548 (1953). 6. W. H. HARTFORD and M. DARRIN. Chem. Rev. 58, 1 (1958); F. and R. H. DuBols. J. Org. Chem. FREEMAN, P. J . CAMERON, 33, 3970 (1968); F. FREEMAN,P. D. MCCARTY,and N. J. YAMACHIKA. J. Am. Chem. Soc. 92, 4621 (1970); K. B. SHARPLESS and A. Y. TERANISHI. J. Org. Chem. 38, 185 (1973); F. W. BACHELOR and U. 0 . CHERIYAN. J. Chem. Soc. Chem. Commun. 195 (1973); Tetrahedron Lett. 3291 (1973); Can. J. Chem. 54, 3383 (1976). K. A. K. L o n , and M. C. R. 7. N. BAILEY,A. CARRINGTON, SYMONS.J. Chem. Soc. 290 (1960). In Oxidation in organic chemistry, Part A. Edited 8. R. STEWART.

17. 18. 19.

20. 21. 22. 23. 24.

25. 26. 27. 28. 29.

by K. B. Wiberg. Academic Press, Ncw York. 1965. p. 35; F. H. WESTHEIMER and A. NOVICK. J . Chem. Phys. 11, 506 (1943). L. D. BURKEand J. F. HEALY.J. Chem. Soc. Dalton Trans. 1091 (1982). D. G. LEEand N. A. NOURELDIN. J. Am. Chem. Soc. 1 0 5 , 3188 (1983). A. STREITWIESER, R. H. JAGOW,R. C. FAHEY,and S. SUZUKI. J. Am. Chem. Soc. 80, 2326 (1958); E. A. HALEVI. Prog. Phys. Org. Chem. 1, 109 (1963); T. H. LOWRY and K. S. RICHARDSON. Mechanism and theory in organic chemistry. 2nd ed. Harper and Row, New York. 1981. pp. 210-212. R. STEWART. The investigation of organic reactions. PrenticeHall, Englewood Cliffs, N.J. 1966. p. 52; F. A. CAREY and R. J. SUNDBERG. Advanced organic chemistry, Part A. Plenum, New York. 1977. pp. 135-139. F. FREEMAN, C. 0 . FUSELIER, C. R. ARMSTEAD, C. E. DALTON, P. A. DAVIDSON, E. M. KARCHESFSKI, D. E. KROCHMAN, M. N. JOHNSON, and N. J. JONES.J. Am. Chem. Soc. 103, 1154 (1981); D. G. LEEand J. R. BROWNRIDGE. J. Am. Chem. Soc. 96, 5517 (1974); K. B. WIBERG,C . J. DEUTSCH,and J. ROCEK.J. Am. Chem. Soc. 95, 3034 (1973). D. F. DETAR.J. Chem. Educ. 44, 191 (1967). D. G. LEE, U. A. SPITZER,J. CLELAND, and M. E. OLSON.Can. J. Chem. 54, 2124 (1976). Q. Rev. Chem. Soc. 15,207 ( 1961); F. BASOLOand J. HALPERN. Mechanisms in inorganic reactions. Wiley, LonR. G. PEARSON. don. 1958. Chapt. 7; K. L. BRIDGES, S. K. MUKHERJEE, and G. GORDON.Inorg. Chem. 11, 2494 (1972); J. P. BIRK.J. Am. Chem. Soc. 91, 3189 (1969); J. H. ESPENSON. J. Am. Chem. SOC.92, 1880 (1970). D. G. LEEandR. STEWART. J. Am. Chem. Soc. 86,3051 (1964). S. W. BENSON.The foundations of chemical kinetics. McGrawHill, New York. 1960. p. 33; K. J. LAIDLER. Chemical Kinetics. McGraw-Hill, New York. 1950. pp. 22-24. and W. P. GRIFFITH. J. Chem. Soc. Chem. ComM. SCHRODER mun. 58 (1979); D. G. LEE, D. T. HALL,and J. H. CLELAND. Can. J. Chem. 50, 3741 (1972); D. G. LEE and J. R. BROWNRIDGE. Can. J. Chem. 51, 2102 (1973). E. J. COREY,R. L. DANHEISER, S. CHANDRASEKARAN, G. E. S. D. LARSEN, P. SIRET,and J. L. GRAS.J. KECK,B. GOPALAN, Am. Chern. Soc. 100, 8034 (1978). R. M. COATES,P. D. SENTER,and W. R. BAKER.J. Org. Chem. 47, 3597 (1982). U.S. A. I. DALTON,JR., H. J. DORAN,and R. D. H. MURRAY. Patent No. 4,225,694 (1980). and A. F. CUNNINGHAM, JR. J. Org. Chem. 48, W. G. DAUBEN 2842 (1983). and C. R. HURLEY. J. Am. Chern. Soc. 74, 5012 R. E. CONNICK (1952); G. NOWOGROCKI and G. TRIDOT.Bull. Soc. Chim. Fr. 684 (1965); D. L. BURKEand J. F. HEALY.J. Chem. Soc. Dalton Trans. 1091 (1982). D. G. LEE and M. VAN DEN ENGH.In Oxidation in organic chemistry, Part B. Edited by W. S. Trahanovsky. Academic Press, New York. 1973. Chapt. 4. and M. C. R. SYMONS.J. Chem. Soc. 284 A. CARRINGTON (1960). I. M. KOLTHOFFand E. B. SANDELL. Textbook of quantitative inorganic analysis. 3rd ed. Macmillan, New York. 1952. p. 337. D. A. SKOOGand D. M. WEST.Fundamentals of analytical chemistry. 2nd ed. Holt, Rinehart and Winston, New York. 1969. pp. 422-423. D. G. LEE and K. C. BROWN.J. Am. Chem. Soc. 104, 5076 (1982).