Homogeneous Catalysis of a Multistep Chain Oxidation

41 Homogeneous Catalysis of a Multistep

Downloaded via UNIV OF CALIFORNIA SANTA BARBARA on July 9, 2018 at 09:04:10 (UTC). See https://pubs.acs.org/sharingguidelines for options on how to legitimately share published articles.

Chain Oxidation J. F. FORD and V. O. YOUNG British Petroleum Co., Ltd., Sunbury-on-Thames, Middlesex, England R. C. PITKETHLY British Petroleum Co., Ltd., Sunbury-on-Thames, Middlesex, England, and The City University, London, England

Co-oxidation of indene and thiophenol in benzene solution is a free-radical chain reaction involving a three-step propa gation cycle. Autocatalysis is associated with decomposition of the primary hydroperoxide product, but the system ex hibits extreme sensitivity to catalysis by impurities, particu larly iron. The powerful catalytic activity of N,N'-di-secbutyl-p-phenylenediamine is attributed on ESR evidence to the production of radicals, probably >NO., and replace ment of the three-step propagation by a faster four-step cycle involving R., RO ., >ΝΟ., and RS. radicals. Added iron complexes produce various effects depending on their composition. Some cause a fast initial reaction followed by a strong retardation, then re-acceleration and final decay as reactants are consumed. Kinetic schemes that demon strate this behavior but are not entirely satisfactory in detail are discussed. 2

/^o-oxidation of indene and thiophenol takes place readily if the reac^ tants i n benzene solution are shaken with oxygen at temperatures i n the range 20° to 40°C. (7). The major primary product has been shown to be irans-2-phenylmercapto-l-indanyl hydroperoxide, I, which re arranges spontaneously to the two racemes of irans-2-phenylsulfinyl-lindanol, II (S), and a tentative reaction scheme involving a three-step radical chain based on the suggestion of Kharasch, Nudenberg, and Mantell (11) was proposed for the formation of I. These three products accounted for 86% of the oxygen absorbed. 207

Mayo; Oxidation of Organic Compounds Advances in Chemistry; American Chemical Society: Washington, DC, 1968.

208

OXIDATION O F ORGANIC COMPOUNDS

Π

Concurrently with our work, Oswald and co-workers published a series of papers that greatly extended the knowledge of the scope of co-oxidation of unsaturated compounds with thiols (15). The results of their work were also consistent with the tentative reaction mechanism of co-oxidation. The spontaneous rearrangement of I has been the subject of a detailed kinetic study (6) which showed that a synchronous hydrogen exchange took place i n a dimeric complex. The rate of rearrangement, although sufficient to reduce the yield of hydroperoxide (typically to 75 weight % ) , is not great enough to affect the kinetics of co-oxidation significantly. However, the rearrange ment is accompanied b y a side reaction, also second-order, which is probably another mode of decomposition of the dimeric complex. This side reaction is catalyzed by iron complexes and is believed to be the chain branching step and source of radicals which gives rise to the autocatalytic progress of co-oxidation i n aromatic solvents. This paper discusses three aspects of our extensive but as yet incom plete studies of the reaction: oxygen uptake i n rigorously purified mate rials, the effects of added iron complexes, and the influence of a wellknown radical capture agent. A n interesting feature of the last, antici pated from the work of Rosenwald (19) on inhibitor sweetening, was the large accelerating effect of N,N'-di-sec-butyl-p-phenylenediamine on the co-oxidation reaction. Attention is drawn to the complex kinetics which can result from delicately balanced catalysis of initiation and termination reactions and changes in the sequence of propagation reactions. Experimental Purification of Reactants and Solvents. Indene and thiophenol of the origin and purity already described (8) were further purified as follows. Indene was percolated under oxygen-free nitrogen through a column containing half its volume of acid-free alumina (obtained from British D r u g Houses). Portions of about 10 volume % of the eluate were discarded at the beginning and end. The heart cut was washed four times with equal volumes of 15 weight % N a O H , with 15% iron-free hydrochloric acid ( B D H electronic grade), and finally with water (dis tilled water percolated through Permutit Biodeminrolit). The material thus obtained after drying over anhydrous analytical reagent grade mag nesium sulfate had a boiling point of 67°C./13 mm., UD 1.5765, freezing point 1.549°, and purity determined by freezing point not less than 99.5%. Examination by gas-liquid chromatography and infrared failed to detect any impurities. Immediately before use, this material was percolated under nitrogen through silica gel which had been freed from iron b y washing with iron-free HC1 until the washings showed no color change u


41.

FORD E T A L .

209

Multistep Chain Oxidation

with ammonium thiocyanate and freed from acid by washing with ironfree water. Use of quartz or borosilicate glass vessels for the final opera tion, providing both were first washed with iron-free HC1 and water, made no significant difference in the co-oxidation. The total iron content of indene thus obtained was about 0.7 p.p.m. compared with about 5 p.p.m. before treatment. Thiophenol was percolated through iron-free silica gel. The total iron content of the thiophenol thus obtained was about 0.5 p.p.m. Benzene (AnalaR grade) was distilled over sodium and percolated through ironfree silica gel. The total iron content of benzene thus obtained was about 0.14 p.p.m. Isooctane was redistilled and similarly treated. Curves A , A ' and Β , Β ' of Figure 1 were duplicate runs on two separate batches of components prepared by this procedure. Anhydrous ferric chloride was used as received from British D r u g Houses. 1,2-Bis(salicylidene amino) ethane iron (III) chloride. [Fe(bissal e n ) C l ] was prepared in the usual way from ferric chloride and 1,2-bis(salicylidene amino)ethane. F o u n d : C , 53.26; H , 4.12; N , 7.44; Fe, 15.6; CI, 9.84. Calculated for C H N 0 C l F e : C , 53.7; H , 3.91; N , 7.83; Fe, 15.6; CI, 9.93 weight % ). This was a high melting (above 3 0 0 ° C ) , nonrecrystallizable brown solid. Ferric acetylacetonate was prepared in the usual way from freshly precipitated ferric hydroxide (m.p., 183.5°C. (d.) (corr). F o u n d : C , 50.8; H , 5.98; Fe, 15.48. Calculated for C H 0 F e : C , 50.9; H , 5.95; Fe, 15.8 weight % ) . 2V,N'-Di-^ec-butyl-p-phenylenediamine ( D S B P D ) was commercially available material ( I C I Topanol M ) redistilled (b.p., 105°C./0.05 mm.). Examination by infrared failed to detect any impuri ties. The first fraction obtained from the distillation contained p-phenylenediamine (2 weight % on total D S B P D ) . Examination of Co-Oxidation Products. Co-oxidation products were identified and hydroperoxide and sulfoxide yields were measured as pre viously described (7, 8). The observed stoichiometric ratio of indene, thiophenol, and oxygen consumed in the whole reaction is 1:1.06:1 (8). Oxygen Uptake-Time Measurements. These were carried out, usu ally at 20° =h 0.1°C. and at 760 ± 5 mm. total pressure (680 mm. partial pressure of oxygen) on 5-ml. samples using the apparatus described by Bolland (2). Reactants dissolved in the appropriate amounts of solvent were introduced into a 15-ml. borosilicate glass flask (previously i n con tact with iron-free H C 1 for 24 hours and washed with iron-free water) cooled i n dry ice-acetone using calibrated pipetes (similarly treated). Co-oxidations were usually carried out at [indene] = [ P h S H ] = 0.15M in benzene. Replacing washed borosilicate glassware with quartz or by plastic ( polyethylene ) made no significant difference in the oxygen-uptake vs. time curves. Iron Estimations. These were carried out by extracting the iron with "iron-free" HC1 (^-Ό 2 p.p.m.), followed by measurement using a square wave polarograph. Reproducibility and accuracy on samples in which known amounts of ferric acetylacetonate had been placed were poor ( ± 5 0 % ) . Wet combustion techniques followed by estimations using the square wave polarograph gave worse results than the procedure described. 1 6

1 4

2

2

1 5

2 1

6

0

0

?


210

OXIDATION O F

ORGANIC COMPOUNDS

Π

Addition of DSBPD during Co-Oxidation at 20°C. A co-oxidation i n the absence of D S B P D (5 m l . solution, 0.15M in thiophenol and indene) was allowed to proceed until 0.08 mole of 0 per mole of P h S H had been absorbed; the reaction mixture was then frozen. A solution of 5 m l . of benzene containing 9 Χ 10" gram of D S B P D was added, and the cooxidation was restarted. The subsequent rate obtained was 0.0086 com pared with an initial rate of 0.015 mole of 0 per mole of P h S H per minute obtained when a similar solution of D S B P D was added at the start of a co-oxidation. 2

4

2

Results and Discussion Co-Oxidation in absence of Added Catalysts or Inhibitors. Figure 1 ( curves Α,Α', Β,Β' ) shows typical oxygen uptake vs. time curves obtained after using the procedures for purifying indene, thiophenol, and benzene and cleaning the reaction vessel as described above. O n the experience of several hundred runs, the reproducibility expected was such that the time for uptake of 0.5 mole of 0 per mole of P h S H would be i n the range 150 to 200 minutes. Despite all efforts, it was estimated that iron 2

1.00

50

Figure 1.

100

TIME ,min.

150

200

250

Co-oxidation of indene and 0.15M thiophenol in benzene at 20°C.

Α,Α', Β,Β': Punfied reagents, no additives C: 0.0075M (PhCOOh D: 0.015M (PhCOOh + 0.0075M PhNMe 2


41.

211


FORD E T A L .

was still present i n the reactants and solvent, giving the curves of Figure 1, to a total concentration in the reaction mixture of about 0.2 p.p.m. If these procedures were not used, curves like those obtained with added iron (Figure 2) were sometimes obtained, and generally, in the presence of added iron complexes, greater variability was found—e.g., curves 375R to 377R in Figure 2. Neither the indene nor the thiophenol, alone, absorbed measurable amounts of oxygen under the conditions of a cooxidation run. i.oo

375R

340R

^ ^ ^ ^ ^ ^ ^ ^

X

.

378R/

^

if

φ /

ο ε cr α 0.5 oc ο

376

/

I90R/

(0 (0

< ζ ω ? δ 352R 50

Figure 2.

10-m

Fe(acac) 352R,1.26 378R,1.48 375R, 5.89 376R, 5.89 377R, 5.89

I0" M J0" M lfr M 10- M 10~ M

s

150 TIME/nin.

200

250

Effect of iron complexes on co-oxidation of indene and thiophenol in benzene at 20°C.

FeCU 189R, 3.14 X 190R, 3.14

100

X J0-5M

X X X X X

5

5

5

5

5

Fe(bissalen)Cl 340R, 3.0 Χ 10-5M

Benzoyl peroxide (Figure 1, curve C ) or even more effectively, benzoyl peroxide plus dimethylaniline (Figure 1, curve D ) , accelerated the reaction to varying degrees. Strong bases altered the course of the


212

OXIDATION O F

ORGANIC COMPOUNDS

Π

reaction by accelerating the direct oxidation of thiophenol to diphenyl disulfide. It was initially assumed as in other autoxidation systems, that autocatalysis arose from first- or second-order radical-producing decomposi tions of the hydroperoxide product of the reaction, and this led to Reac tion Scheme 1 (shown below). Initiation Indene + 0

—» radicals

(1)

R O O H (I)

->

RO* + OH*

(2)

2ROOH

-»

R 0 * + RO* + H 0

(3)

->

PhS* + R 0 H

(4)

2

or some unknown source

2

2

Propagation R0

2

+PhSH

SPh

2

SPh

Termination Bimolecular disproportionation and/or recombination of radicals

(7)

Reaction Scheme 1 The greater speed of the co-oxidation compared with autoxidation of the indene alone was accounted for by assuming that Reactions 4 and 5 are faster than addition of R 0 * to indene. Analyzing the kinetics of this system ( 18 ) showed that linear combi nation of the three initiating reactions could account for long accelerating periods but could not explain a constant or slightly falling rate which was sometimes observed in the early stages of co-oxidation. The kinetic analysis also suggested that the slow rise and general shape of the oxygen-uptake curves could be accounted for if the radical buildup were slow and nonstationary. This explanation was rejected after several ex periments in which a rapidly progressing co-oxidation was quenched in liquid air, and the mixture was allowed to warm to 20°C. under nitrogen 2


41.

FORD E T A L .


213

and stand for a time. After subsequent refreezing, replacement of nitro gen by oxygen, and rewarming, the reaction proceeded at exactly the rate observed just before the interruption. Thus, unless the free radicals were startlingly long-lived, the initiation must come from stable products, and the reaction is in a pseudo-steady state—i.e., the main propagation chain responds rapidly to changes in initiation and termination rates. The influence of the products of the reaction on the kinetics was then examined. First, the stable products, isolated from the reaction mixture after rearrangement of hydroperoxide was complete (S), were added (5 weight % on P h S H ) in separate experiments to the indenethiophenol reaction mixture. The racemic sulfoxides, II, the correspond ing frans-hydroxysulfide, a ketosulfoxide concentrate, and freshly recrystallized diphenyl disulfide, had no significant effect on the oxygen uptake vs. time curves. Second, portions of the whole co-oxidation mixture after oxygen absorption ceased, and containing about 0.11 mole per liter of hydroperoxide, I, when added to a fresh reaction mixture greatly in creased the oxidation rate. However, varying the amounts of added fresh co-oxidation products showed a complex and rather irreproducible con centration dependence. The effect of partly decayed co-oxidation prod ucts was also unexpectedly complex, and completely decayed products containing no hydroperoxide could vigorously accelerate co-oxidation. This confirmed that simple first- or second-order radical-producing de compositions of peroxide could not adequately explain the catalytic behavior of the products, although the peroxide undoubtedly played an essential role in the autocatalysis of the main reaction itself. Peroxide in solution is necessary to produce fast co-oxidation. W h e n isooctane was used as solvent instead of benzene, a slow and linear oxygen uptake was observed (0.5 mole per mole of indene in 100 hours at 20°C. ). This behavior was the result of low solubility in the paraffin of the hydro peroxide, I, which precipitated out and rapidly rearranged to inactive sulfoxides. W e conclude that although Scheme 1 represents a basic mechanism for co-oxidation of indene and thiophenol, the system is so sensitive to the effects of catalysts and inhibitors that interpreting its behavior neces sitates studying the initiation and termination mechanisms. Co-Oxidation in Presence of Added Ν,Ν'-di-sec-butyl-^-phenylenediamine. Figure 3 shows typical oxygen uptake curves obtained if D S B P D is added at concentrations around 10" to 10" M to the indene-thiophenol reaction mixture. The reproducibility was good. The products after oxygen absorption was complete and the hydroperoxide had been allowed to decay were substantially the same as those obtained in the absence of D S B P D . Uptake of oxygen was extremely slow when either of the main reactants was omitted. 2

4


214

OXIDATION O F O R G A N I C C O M P O U N D S

Π

1.00

TIME,min.

Effect of Ν,Ν'-di-sec-butyl-p-phenylenediamine on co-oxidation of indene and 0.15M thiophenol in benzene at 20°C.

Figure 3.

A: no additive

B: 7.5 Χ I0-5M DSBPD

C: 7.5 X I0- M preoxidized DSBPD D: 7.5 Χ 10-3M DSBPD 5

It is reasonably well established that aromatic amines act as inhibi tors of hydrocarbon autoxidation by capturing the chain-propagating radical ( R 0 * ) to form, eventually, species which do not propagate the chain. Thus, i n competition for R 0 * with a hydrocarbon, the amine is successful. In the co-oxidation reaction, since P h S H is known to be an efficient radical scavenger or polymerizing chain transfer agent, R S H might be expected to compete successfully with the amine for R 0 * . This would explain a lack of inhibiting action of D S B P D i n the co-oxidation but not the large accelerations it can produce, nor the high initial rates observed with no significant induction period. Rosenwald ( 19 ) attributed the acceleration of the reactions involved in inhibitor sweetening to the reaction of an unidentified oxidation prod uct of the amine with thiol to produce mercaptyl radicals. However, the oxidation of neat D S B P D is slow at 20°C. (3.3 Χ 10" mole of 0 per mole of amine per second ) and partial pre-oxidation of the D S B P D until 0.26 mole of 0 was absorbed per mole of amine d i d not enhance the activity of the amine (compare curves Β and C i n Figure 3 ) . Further more, allowing the co-oxidation to proceed for a time before adding the 2

2

2

δ

2


2

41.

FORD ET AL.

215


D S B P D d i d not significantly alter the accelerated rate of co-oxidation. It must, therefore, be concluded that the initiation arises from interaction of the amine with a reactant and not with a product such as hydroper oxide. In addition, Rosenwalds "oxidized state" of the amine cannot be among the stable oxidation products of the amine. The reactions of D S B P D were therefore examined. D S B P D d i d not appear to react with thiophenol under nitrogen, but it d i d react slowly with oxygen. Pure D S B P D showed no E S R spectrum, but bubbling oxygen through the sample tube was followed by the appearance of a red color and a three-line spectrum (Figure 4 ) , with g = 2.003 and splitting a = 11.9 gauss and partly resolved fine structure (a = ca., 2.3 gauss). After the oxygen was pumped oflE, the three-line spectrum faded. N

H

11-9 G A U S S 11-9 G A U S S

Figure 4.

N,N'-Di-seC'butyl-p-phenylenediamine

plus oxygen at

20°C.

This behavior is not consistent with the formation of Wurster ions, which have been postulated as intermediates and i n fact demonstrated to occur i n polar solvents (3) but indicates the formation of a nitroxyl radical intermediate in the presence of oxygen and its decay when the oxygen is removed. [Sullivan (23) gives a number of references to nitroxyl radicals.] The possibility that R N * radicals are formed as a 2


216

OXIDATION O F

ORGANIC COMPOUNDS

II

result of hydrogen abstraction by radicals and the problem of distinguish ing the E S R spectra of R N * and R N O ' radicals are discussed by Coppinger and Swalen (5). They favor decomposition mechanisms involving oxidation-reduction chains between > N O ' and > N O H . Hydrogen ab straction reactions involving diphenylnitroxyl- and diphenylhydroxylamine have been postulated to account for catalysis of benzoyl peroxide decomposition (4). Admitting thiophenol to an oxidizing sample of D S B P D immediately quenched the three-line E S R spectrum. Also D S B P D containing thio phenol absorbed oxygen much more rapidly than the amine itself and formed diphenyl disulfide and water, but no color appeared. However, thiophenol does not discharge the color of an oxidized sample of the amine. Thus, although the nitroxyl radical may be red, the color of the oxidized amine is mostly associated with the products of further oxidation. It is evident that the > N O * radical from D S B P D reacts readily with thio phenol, presumably abstracting hydrogen and producing PhS* radicals. Other more stable > N O * radicals have been reported to behave similarly (9). Baird and Thomas ( I ) have shown that radical initiators—e.g., A Z D N in the presence of oxygen—produce > N O ' radicals from primary and secondary amines and Thomas (24) proposed a mechanism for amine attack by R 0 ' radicals: 2

2

2

> N H + R0 * - » >N + ROOH 2

>N* + R 0 * -> > N O ' + RO* 2

A n alternative mechanism for nitroxyl radical production analogous to one of the higher temperature modes of R 0 * destruction is: 2

>N* + 0

2

- » >NOO*

2 >NOO* -> 2 >NO* + 0

2

These reactions are probably fast, and a plausible modification of Reaction Scheme 1 gives Scheme 2, which fits the observed facts. If the > N ' radicals are more active than > N O * , they could play a more important role in the chain propagation; however, the rate of reaction of P h N * with thiols is relatively slow and even at 100°C. is incomplete after more than 100 hours (26). Only the radical which has been identified is included i n Scheme 2. 2

Initiation at Start

0

2

+ DSBPD

Later R O O H Propagation

- » >NO*

(8)

—> radicals

(2)

2 ROOH

—> radicals

(3)

PhS' + indene

- » R'

(5)

R* + 0NOH

- » >NO* + R O O H

(9)

>NO* + PhSH

- » > N O H + PhS'

(10)

2

Termination

217

Disproportionation and/or combination of radicals

(Τ)

DSPD + R 0 *

- » inert products

(ID

>ΝΟ'

—» inert products

(12)

2

Reaction Scheme 2 In Reaction Scheme 2 the acceleration of co-oxidation by D S B P D arises from > N O ' radical production by direct interaction of the amine with oxygen in Reaction 8, together with the introduction of a pair of fast steps (Reactions 9 and 10) which bypass the rate-determining step of the main propagation chain (Reaction 4 of Scheme 1). The main propagation chain has now become a four-step chain consisting of Reac tions 5, 6, 9, and 10. Consumption of the D S B P D results from further oxidation of the > N O ' radical (Reaction 12) but may also occur by hydrogen abstraction from the alkyl group leading to imine formation (5) or by dispropor tionation reactions yielding quinonoid structures (9). The existence of this reaction system may therefore be transitory. However, the main chain is long, and the effect of 7.5 X 10" M D S B P D lasts almost through out the oxidation of 0.15M indene-thiophenol (see Figure 3 ) . Co-Oxidation in Presence of Added Iron Complexes. Our study of the influence of added iron on the co-oxidation is not yet complete, and we have encountered experimental difficulties (not yet overcome) i n recognizing changes i n the coordination shells or complexing species and in measuring concentrations of F e and F e at p.p.m. concentrations i n hydrocarbon solvents during the co-oxidation. However, several inter esting observations have been made, and several possible reaction schemes have been examined. A wide range of iron compounds catalyze the indene-thiophenol co-oxidation reaction—e.g., F e C l and a variety of complexes formed by chelation. Differences i n their behavior are believed to be correlated approximately with the ease of displacement of ligands by thiols or with their susceptibility to attack by peroxide, but the lack of reproducibility which has made much of this work extremely tedious makes quantitative deductions difficult. Figure 2 shows typical oxygen-uptake curves obtained after adding ferric complexes or ferric chloride (at about 10" M) to indene-thiophenol reaction mixtures in benzene solution. The three consecutive runs (375377R) made under apparently identical conditions reasonably represent the variability found. 3

3+

2+

3

4


218

OXIDATION O F

ORGANIC COMPOUNDS

II

The products of co-oxidation examined after oxygen absorption has ceased and after the hydroperoxide had rearranged were substantially the same as those obtained in absence of added iron, although some changes in the yields of minor products resulted from the effects of the iron complexes on the decomposition of the hydroperoxide. The reactions when ferric chloride or ferric acetylacetonate is present occur in three stages: (1) an initial fast reaction decreasing in rate, (2) an arrest during which the oxidation rate may be very low, and (3) a "sigmoid" phase i n which an autocatalytic reaction is finally overtaken by reactant consumption. W i t h the ferric (bissalen) chloride complex, the reaction was accelerated but Stages 1 and 2 did not appear (curve 340R, Figure 2). In examining the kinetics of this type of chain reaction (18), it was found that the major characteristics of all the curves of Figure 2 could be reproduced by assuming: ( a ) Radicals were produced in two reactions, the rate of one being independent of the extent of reaction and that of the other being propor tional to the square of the extent—e.g., the latter might be a second-order decomposition of the main peroxide product and might be catalyzed by iron. (b) Radicals were removed in two ways, by bimolecular termination processes and by a reaction, whose rate is proportional to the extent of reaction. Such a system was delicately balanced, and small changes in the arbitrary constants could greatly alter the shapes of the reaction curves. However, before claiming that this system represents the reactions which are occurring, a plausible explanation must be found for the termination reaction, which had a rate proportional to the extent of reaction. T w o possible assumptions have been considered, but neither fully stands up to experimental test, and it is evident that the system is one of subtle balance and complexity. It was first assumed that an inhibitor was produced as a by-product. However, the accuracy with which the initiation rates had to be balanced seemed to demand too great a coincidence if the inhibitor production occurred in a side reaction independent of the initiation reactions. Thus, although possible products which were demonstrably powerful inhibitors of the co-oxidation reaction—e.g., thiolsulfinates—could be postulated, this mechanism was unsatisfactory and d i d not fully explain the role of iron. Furthermore, diphenyl disulfide, which is a precursor of thiolsulfinate in the presence of hydroperoxide, did not inhibit the reaction. The possibility that the negative term represented a diminishing contribution from a radical-producing process was then studied. Careful examination of the oxygen-uptake curves with ferric acetyl acetonate or ferric chloride present, failed to reveal an induction period


41.

FORD E T A L .

219


or initial accelerating phase which, if it occurred, must have lasted less than a minute. In view of the rigorous purification procedures used, the occurrence of a high initial rate in Stage 1 does not support the proposi tion that the iron accelerates the initial stages of co-oxidation by the usual redox interactions with a hydroperoxide product, whether this arises from direct oxidation of indene (20) or from the co-oxidation reaction. Furthermore, ferric acetylacetonate is not an effective catalyst for the autoxidation of indene, as we have found, nor of 1-octene (10), and a concentration of 1.8 X 10" M gives only a fivefold increase in the rate of decomposition of hydroperoxide I to products other than II at 20°C. 4

In order to produce a high initial rate we suggest that the iron com plex must produce radicals by attacking a reactant, and the thiol is the most likely one. This proposal is supported by the recent demonstration by Wallace (25) that ferric octanoate readily reacts with thiols at ambient temperature to give R S ' radicals which are effectively captured by an olefin, provided the ratio of thiol to iron concentrations is not greater than 10. Ferric ions and complexes in aqueous media react readily with cysteine (13, 21) or thioglycollie acid (14, 22) to form purple complexes which rapidly change to the ferrous form with accompanying formation of disulfides. It was not unreasonable, therefore, to assume that such a reaction provided rapid production of RS* radicals in the initial stages of co-oxidation and that this rate should decrease as reduction to ferrous occurred. Eventually reoxidation of ferrous to ferric by peroxide would assert itself, and the rate could rise again. The system including all these assumptions has not yet been calcu lated, but the slightly simplified Reaction Scheme 3 has been examined. Although it was also found inadequate, it has some interesting kinetic properties. Initiation

Propagation

2R0 H

-> RO* + R 0 + H 0

Fe

3+

+ PhSH

- » Fe

Fe

2+

+ R O O H -> F e

2

2

2

2+

+ PhS + H

+

3+

+ RO' + O H -

(13)

k

lt

(14)

RO* + PhSH

- » PhS* + R O H

Fast

R 0 * + PhSH

—» PhS' + R O O H

k

(4)

h

(5)

- » R0 *

h

(6)

—» inert products

k

2

PhS* + indene -> K R' + 0 Termination

(3)

*3

2 R0 * 2

2

2

4

7

Reaction Scheme 3


(7)

220

OXIDATION

OF

ORGANIC

COMPOUNDS

II

Assuming a pseudo-steady state in which initiation and termination rates may change but remain equal and that Reaction 6 is fast compared with 4 and 5: Rate of radical production (RCV)

r = k

7

2

and

N O ' radicals which are capable of reacting with thiophenol to produce PhS' radicals, thus initiating the reaction, and by replacing the rate-determining propa gation step by a pair of faster reactions involving > N O * (or > N ) radi cals, thus converting the three-step chain into a four-step chain. Iron complexes have profound effects, depending on their stabilities. Many markedly accelerate the initial stages of the co-oxidation reaction but retard and even stop the reaction i n the middle stages. The kinetic features can be accounted for largely by a balance among several initia tion and termination reactions, but discrepancies suggest that changes occur in the composition of the complexes as a result of ligand exchange or displacement, attack on ligands b y free radicals, and changes i n oxida tion state of the metal during oxidation, and that these are accompanied by changes i n reactivity and therefore i n catalytic activity. Systematic study of the stability of metal complexes i n nonaqueous systems and their reactivity with thiols, free radicals, etc., is required for a fuller understanding of their complex effects on the co-oxidation reaction. Acknowledgment The authors thank the British Petroleum Co., L t d . , for permission to publish this paper. They are indebted to J. C . Blackburn for the E S R measurements and interpretation and to R. C . Palmer for developing the system of equations and for programming and arranging the computa tions b y digital computer. Literature Cited (1) Baird, J. C., Thomas, J. R., J. Chem. Phys. 35, 1507 (1961). (2) Bolland, J. L., Proc. Roy. Soc. (London) 186A, 218 (1946).


224

OXIDATION OF ORGANIC COMPOUNDS—II

(3) Boozer, C. E., Hammond, G. S.,J.Am. Chem. Soc. 26, 3861 (1954). (4) Chalfont, G. R., Hey, D. H., Liang, C. S.Y.,Perkins, M. J., Chem. Com mun. No. 8, 367 (1967). (5) Coppinger, G.M.,Swalen, J. D., J. Am. Chem. Soc. 83, 4900 (1961). (6) Ford, J. F., Pitkethly, R. C., Young, V. O., National Bureau of Mines Conference on Low Temperature Oxidation, Bartlesville, Okla., May 1961. (7) Ford, J. F., Pitkethly, R.C.,Young, V. O., "Preprints of Papers," Division of Petroleum Chemistry, ACS Meeting, Miami, Fla., 2 (1), 111 (April 1957). (8) Ford, J. F., Pitkethly, R.C.,Young, V. O., Tetrahedron 4, 325 (1958). (9) Forrester, A. R., Thomson, R. H., Nature 203, 74 (1964). (10) Indictor, N., Jochsberger, T., J. Org. Chem. 31, 4271 (1966). (11) Kharasch, M. S., Nudenberg, W., Mantell, G. J., J. Org. Chem. 16, 524 (1951). (12) Mendelsohn, M., Arnett, E. M., Freiser, H., J. Phys. Chem. 64, 660 (1960). (13) Michaelis, L., Barron, E. S. G., J. Biol. Chem. 83, 191, 367 (1929); 84, 777 (1930). (14) Michaelis, L., Schubert, M., J. Am. Chem. Soc. 52, 4418 (1930). (15) Oswald, Α. Α., Wallace, T. J.; in "Chemistry of Organic Sulfur Com pounds," Vol. 2, Chap. 8, p. 217, N. Kharasch and C. Y. Meyers, eds., Pergamon Press, 1966. (16) Patmore, E. L., Ph.D. thesis, University of Connecticut, 1963; University Microfilms Inc., Ann Arbor, Mich., 1966. (17) Patmore, E. L., Gritter, R.J.,Proc. Chem. Soc. 1962, 328. (18) Pitkethly, R.C.,unpublished manuscript. (19) Rosenwald, R. H., Petrol. Processing 6 (9), 969 (1951); 11 (10), 91 (1956). (20) Russell, G. Α., J. Am. Chem. Soc. 79, 3871 (1957). (21) Schubert,M.,J.Am. Chem. Soc. 53, 3851 (1931). (22) Ibid. 54, 4077 (1932). (23) Sullivan, A. B., J. Org. Chem. 31, 2811 (1966). (24) Thomas, J. R., J. Am. Chem. Soc. 82, 5955 (1960). (25) Wallace, T. J., J. Org. Chem. 21, 3071 (1966). (26) Wallace, T. J., Mahon, J. J., Kelliher, J.M.,Nature 206, 709 (1965). RECEIVED November 13, 1967.

Discussion H . Berger: D o you think that the fact that strong bases do not lead to co-oxidation but to disulfide, is an indication that no RS · radicals are involved i n this side reaction? R. C . Pitkethly: Yes, our experience and that of others ( I , 2, 3) would indicate that the oxidation of disulfides i n alkaline media involves


41.

FORD E T A L .


225

formation and reaction of metal complexes and that free RS · radicals are not formed i n any appreciable concentration. Literature Cited (1) Hill, J., McAuley, Α., Chemical Society Autumn Meeting, Sept. 20, 21, 1967, Paper B4. (2) Overberger, C. G., Burg, K. H., Baly, W. H., J. Am. Chem. Soc. 87, 4125 (1965). (3) Swan, C. J., Trimm, D. L., Chem. Ind. 1967, 1363.


Homogeneous Catalysis of a Multistep Chain Oxidation

Recommend Documents