Chapter 5 The Oxidation of Organic Compounds by Non-metallic Anions

Chapter 5

The Oxidation of Organic Compounds by Non-metallic Anions G . J. B U I S T

1. Kinetics of oxidations by periodate 1.1

INTRODUCTION

Periodate is well known as a specific oxidant in organic chemistry, and it has been used extensively in structural studies of carbohydrates (see reviews by Jackson’, Babbitt', and Dryhurst3). Its particular value is that under suitable conditions of pH and temperature its action is largely restricted to the oxidation of certain types of compounds (those classified under (a) below). For example, 1,2-diols and 1,Zdiketones are readily oxidised, whereas most alcohols, aldehydes, and ketones are resistant. However, the action of periodate is not nearly so specific as was thought at one time, and there is a wide range of compounds which it is capable of oxidising (Sklarz4). Periodate oxidations of organic compounds may be classified as follows: (u) oxidations accompanied by carbon-carbon bond fission (at a suitable pH most of these proceed rapidly at room temperature or below), (b) other types of oxidations proceeding at a measurable rate at room temperature, and (c) oxidations whose rate is only appreciable at higher temperatures, or in the presence of ultraviolet radiation. Oxidations of type (a) are often referred to as “Malapradian” after the discoverer of the fission of 1,Zdiols by periodate (Malaprade6, 1928). Other compounds whose oxidations are of this type are l,Zdiketones, l,Zhydroxyketones, 1,2-aminoalcohols, 1,Zdiamines, uhydroxy acids, and a-keto acids. Most kinetic studies of periodate oxidations have concerned the reaction with 1,2-diols, but some work on other oxidations of type (u) has been reported, and will be reviewed in sections 1.4 and 1.5. The very few kinetic studies of oxidations of type (b) will be reviewed in sections 1.6 and 1.7. No kinetic studies of type (c) oxidations have been reported. Bunton’ has reviewed the mechanisms of periodate oxidations. In most of its reactions periodate is a two-electron oxidant, and it is reduced to iodate. Periodate oxidations are normally carried out in aqueous solution, but organic or mixed aqueous-organic solvents can be used (Qureshi and Sklarz’) and in section 1.8 the very limited amount of kinetic data available for mixed solvents is reviewed. References pp. 489-492

436

OXIDATION OF ORGANIC COMPOUNDS

1.2

P E R I O D A T E SPECIES I N A Q U E O U S S O L U T I O N

Periodic acid has the following pK values at 25 "C: pKl = 1.64, pK2 = 8.31, pK3 = 12.2 (Buist et al.'). The acid itself is present exclusively as H5106, but the monoanion H,IO; is largely dehydrated to IO;, and the dianion undergoes dimerisation, uiz.

At 25 "C KD = 40 (Crouthamel et aZ.,9,10)and K, = 141 1.mole-' at an ionic strength of 0.1 (Buist et al.'). The measured first and second dissociation constants, R, and differ from the constants Kl and K, defined above, on account of the formation of 10,

x2,

(the latter equation assumes solutions sufficiently dilute to make dimerisation negligible). The constant K, is markedly temperature-dependent, whereas R , shows a small temperature dependence (Buist et aL8).

1.3

K I N E T I C S A N D MECHANISM O F O X I D A T I O N O F

1,2-DIOLS

The stoichiometry of the oxidation of 1,Zdiols is R'RZC(OH).C(OH)R3R4+10~= R'R2C=O+R3R4C-0+IO; +H,O In this equation, as well as in later stoichiometric equations, the formulation of periodate as 10, is used for convenience, and does not imply that the latter ion is necessarily the reactive periodate species.

1.3.1 Evidence for the formation of diol-periodate esters Diol-periodate esters are recognized as intermediates in the oxidation of 1,2-

1

O X I D A T I O N S BY P E R I O D A T E

437

diols by periodate. Their formation is well substantiated by kinetic evidence (see following sections) and by other evidence summarised here. Criegee’ compared the oxidations of 1,Zdiols by periodate and lead tetraacetate, and suggested that both oxidations proceed via cyclic ester intermediates. However, lead tetraacetate can oxidise 1,Zdiols in which the formation of a cyclic ester is clearly impossible (e.g. trans-9,10-decalindiol),and it is evident that an alternative reaction pathway exists for lead tetraacetate oxidations, whereas for periodate all the evidence suggests that a cyclic ester is a necessary intermediate. In addition to trans-9,lOdecalindiol, several 1,2-diols are known which are inert to periodate, and in all cases it is clear that steric effects make the formation of a cyclic ester very difficult or impossible. The formation of diol-periodate esters is supported by physical evidence. The addition of ethane-l,2-diol to periodate solutions causes an initial rapid change in the uv absorption spectrum, followed by a slower change as the oxidation proceeds and 10; is formed. Similar results are observed for other 1,Zdiols except for highly substituted diols such as pinacol (Buist et al.”). Buist and Bunton’ have shown that the cyclic periodate esters formed in alkaline solution from 1,2-diols and periodate can be detected by NMR. The initial fall in pH which occurs in the oxidations of ethane-1,2-diol and lightly substituted diols is also attributed to ester formation (Mala~rade’~, Buist and Bunton’ ’). Cyclic triesters, similar to the cyclic diesters formed from 1,2-diols, are formed from cyclic compounds containing the cis-1,2,3-triol system (Barker and Shawl6, Dijkstra 17) and from 1,2-0-isopropylidene-cr-~-glucofuranose. In the latter case the presence of the triester has been demonstrated by NMR (Perlin and van Rudloff 18). Monoesters of periodic acid have not been detected in any system, but they are postulated as intermediates in the formation of cyclic diesters from 1,Zdiols (section 1.3.5).

1.3.2 Kinetics of the oxidation of ethane-1,2-diol Two types of kinetics have been observed in the oxidation of 1,2-diols by periodate. The first of these, exemplified by ethane-l,Zdiol, is of a mixed-order type, and is consistent with the reversible formation of the cyclic diester as an intermediate. The second type, exemplified by pinacol, is second-order kinetics, the reaction showing general acid-base catalysis and a complex dependence of rate on pH. The first kinetic investion of the ethane-1,2-diol reaction was carried out by Price and Knelllg. However, these workers did not detect the typical mixed-order kinetics which were later established by Duke”, and investigated over a wide range of conditions by Buist and Bunton’ ’. With excess diol Duke showed that the reaction is first-order with respect to periodate, and that the first-order rate coefficientk’ is a function of diol concentration, viz. References pp. 489-492

438

OXIDATION OF ORGANIC COMPOUNDS

where D represents ethanediol and Kis a constant. Duke explained this relationship in terms of the rapid reversible formation of an intermediate, followed by the slow decomposition of the latter, uiz. (CHzOH)2+IO;

.+Intermediate k K

--f

2 CH20+IO; +HzO

This reaction sequence leads to equation (1) for k', the first-order rate coefficient with respect to t o t d periodate (i.e. free periodate plus periodate combined in the intermediate), uiz. k'

=

kK[D]/(l

+K[D])

(1)

i.e.

l/k' = l/kK[D] + l/k

(2)

By plotting l/k' against l/[D] Duke evaluated the equilibrium constant K and the rate coefficient k for decomposition of the intermediate. Buist and BuntonI5 showed that equation (1) holds for the ethane-1,Zdiol reaction over the range pH 1-9 at 0 "C, and that a similar equation holds when periodate is in excess. Both K and k are pH-dependent, as shown in Fig. 1. The initial drop in pH found in unbuffered solutions (previous section) shows that the intermediate is a stronger acid than periodic acid, in qualitative agreement with the general increase of K with increase of pH. The plateau in the region pH 3-5 for the variation of k with pH suggests that the monoanion of the intermediate is the decomposing species. These conclusions were supported by the quantitative treatment of the pH dependence based on the scheme HJ06

+ Per- Kz+ HJOiEl

products where Per- represents 1 0; + H,IO,,

and H,C, HzC-, and HC2- represent the

I

OXIDATIONS BY PERIODATE

439

undissociated intermediate, its monoanion, and its dianion respectively. The treatment showed that k" and k= are negligible, i.e. H,C- is the only species whose decomposition is appreciable (at least over the pH range studied). Table 1 includes the values of the equilibrium constants and the rate coefficient k- found for the ethane-l72-diolreaction at 0 "C.The first and second dissociation constants, K ; and K i , of the intermediate can be compared with those of periodic acid at O"C, uiz. 4.0 x lou3and 7.9 x

3.0

Y

0 J

o

-

:

2

Fig. 1. Variation with pH of the rate coefficient for the decomposition of the ethanediol-periodate intermediate, and of the equilibrium constant for its formation, at 0 "C.

The salt effect on the reaction is simply a secondary effect on the various equilibria, and there is no detectable kinetic salt effect on the decomposition of H2C-. The intermediate H,C is considered to be the cyclic diester referred to in the previous section. In order to explain why the monoanion H,C- is the only species of diester whose decomposition is appreciable, Buist and Bunton' put forward as a hypothesis that hydrated and dehydrated forms of the ester exist, (I) and (IV), and that the latter and its anion (V) are the only species capable of decomposing to the oxidation products, viz. References pp. 489-492

440

O X I D A T I O N OF O R G A N I C COMPOUNDS

’

H,C-O

/I\

I OH OH

-

I

H2C-~’~\OH

I

-

I 0-

0-

m

I

I

Products

Products

The most favourable range for existence of the dehydrated form is pH 2-6 when it would be present as its anion (V). The hydrated form is expected to be a weaker acid, so below pH 2 the equilibrium between the two forms must be displaced towards the hydrated form. In alkaline solution the hydrated dianion (111) must be the dominant species because its dehydration is impossible. Thus the predicted range of maximum concentration of the dehydrated form coincides with the range in which the observed rate is at its maximum, furthermore the impossibility of forming a dehydrated dianion is consistent with the steady decrease of rate observed for the range pH > 7. In strongly acid solution the decomposition of the dehydrated ester (IV) is expected to be appreciable, but in practice it is likely that kinetic measurements cannot be extended to a sufficiently low pH. Hence, in the pH range normally studied, the dehydrated anion (V) is considered to be the sole entity capable of decomposing to the oxidation products, uiz. H2C=0

H2C-O

+

0

4- 10;

H2C-0

p.

The internal rearrangement which (V) must undergo to give the products is consistent with the negligible kinetic salt effect. Taylor” pointed out that the kinetics observed for the oxidation of ethanediol are equally consistent with a mechanism involving the direct formation of the products, and an inactive diol-periodate ester, C , in equilibrium with the reactants, uiz.

D +Per + Products

11

C

1

OXIDATIONS B Y PERIODATE

441

Provided that the ester C is in equilibrium with the reactants this mechanism is kinetically indistinguishable from that discussed above. However, with certain substituted ethanediols the rate of formation of the ester is slow, and under suitable conditions the ester is not in equilibrium with the reactants. The two mechanisms then become kinetically distinguishable. A study of the oxidation of 2-methylbutane-2,3-diol (see below) showed that the reaction must proceed via an intermediate, and that the mechanism originally suggested by DukeZois at least in some, and probably in all cases, correct. From the results of Taylor" over the temperature range 0"-25 "C,the variation of the rate coefficient k- for decomposition of the ester monoanion, HzC-, is represented by

k-

=

5.7 x loi7 exp (-23,90O/RT)

The equilibrium constant for formation of the ester decreases with increase of temperature, and for the reaction D+Per-

+ H2C-

AH = -16.4 kcal.mole-' (Taylor2'). Note that Per- and HzC- represent all periodate monoanions and all ester monoanions, irrespective of their degree of hydration.

1.3.3 Reaction of other ly2-diols showing the ethane-Iy2-diol type of kinetics The kinetics of the oxidation of methyl substituted ethane-1,2-diols are generally similar to those observed for ethane-1,2-diol itself, with the exception of the oxidation of pinacol (see section 1.3.6). However, under certain conditions the rate of formation of the diester intermediate is comparable with its rate of decomposition to the reaction products, i.e. the ester is no longer in equilibrium with the reactants. The formation of the ester is buffer-catalysed (section 1.3.6) so equilibrium conditions are more likely to be attained when the reaction is carried out in a buffer. One phenyl substituted 1,2-diol has been studied in detail, viz. phenylethane1,2-diol, and the kinetics of its oxidation are similar to the ethane-1,2-diol kinetics (Buist and Lewis"). The kinetics of the oxidation of cis- and trans-cyclohexane1,Zdiols show deviations from equilibrium over much of the pH range studied, but Buist et aLZ3found it possible to determine the constants shown in Table 1. The pH dependence of the rate coefficients and equilibrium constants for all the 1,Zdiols listed in Table 1 is of exactly the same form as that found for ethane1,2-diol. Fromthequantitative treatment of the pH dependence Buist et al.' 5 ~ 2 2 - z 4 ~ References pp. 489-492

442

O XIDAT IO N O F O R G A N I C C OMPOU N D S

TABLE 1 R A T E C OE FFI C I ENT S A N D E Q U I L I B R I U M C O N S T A N T S FOR THE O X I D A T I O N O F BY P E R I O D A T E A T 0

"c

K"

K-

K"

10ZK1

1,2-DIOLS

1 0 7 ~ ' ~ 103k(sec-')

10 39 270 19 16 60 5.6

189 500 68 101

274

2700 4100

8ooO

370 3 60 940 4400 1000

400 ~~

*

7.4 5.0 3.4 2.6 2.6 1.4 19.7 3.1 0.8

1.1 0.65 0.42 0.50 0.28 0.19 1.3 0.85 0.48

4.57 12.0 30.0 18.2 61 24.4 66 16.5 330

~

The data for this diol are for a temperature of 1 "C.

calculated the equilibrium constants, and the rate coefficients for decomposition of the diester monoanion, given in Table 1. The symbols refer to the reaction scheme on p. 438. The effect of substitution on the equilibrium constants KO, K - , and K = is discussed in the next section. The sequence of values observed for the first and second dissociation constants of the diester (K; and K l ) are consistent with the normal inductive effects of the methyl and phenyl groups. The general increase of the rate coefficient k - with methyl substitution is presumably due to hyperconjugation between the methyl groups and the partially formed carbonyl groups in the transition state for the decomposition of (V). The corresponding mesomeric effect for the phenyl group is expected to be more powerful, in accordance with the greater k- value observed for phenyl-ethane-l,2-diol compared with propane1,2-diol. The very high k - value for cis-cyclohexane-l,2-diolmay be due to the release of steric strain in the fused ring system (Buist et ~ 1 . ' ~ ) .

1.3.4 Structure of the diol-periodate cyclic diester Because the undissociated diester and its monoanion may exist in hydrated and dehydrated forms (section 1.3.2), discussion of the structure of the diester has been focussed on its dianion which can only exist in the hydrated form (111). By analogy with periodates of known structure, the periodate moiety of (111) is assumed to have an octahedral arrangement of oxygen atoms. Molecular models show that the 5-membered ring in (111) is likely to be puckered, as illustrated in Fig. 2, and that two of the possible positions for substituents, H and H' (hindered positions) are comparatively close to two of the periodate oxygen atoms, whereas the other two positions, F and F', (free positions) are not close to any periodate

1

OXIDATIONS BY PERIODATE

H

443

P

F

H’

Fig. 2. Model of the cyclic diol-periodate ester.

oxygen atoms (Buist et ~ 1 . ’ ~ )This . model helps to explain the sequence of values of K’ in Table 1. The inductive effect of methyl substitution increases the stability of the diester (with respect to the reactants) but a methyl group in a hindered position introduces steric compression and a reduction in stability. For propane1,Zdiol and (f)-butane-2,3-diol the methyl groups can be in free positions, but the remaining methyl-substituted diols in Table 1 must have one methyl group in a hindered position. It is precisely the latter diols which have reduced values of K=. Similar considerations explain the greater K = value observed for transcompared with cis-cyclohexane-1,Zdiol (Buist et ~ 1 . ’ ~ ) . Buist and B ~ n t o n ’have ~ obtained the NMR spectra of the diesters formed from several 1,2-diols. The chemical shifts and the coupling constants confirm the structure illustrated in Fig. 2.

1.3.5 Kinetics of formation of the diol-periodate diester In the range p H 9-11 the rate of formation of the diester is slow enough at 0 “C to be followed by conventional spectrophotometry. The kinetics are those of a reversible reaction, second-order in the forward direction and first-order in the reverse direction, in agreement with the reaction scheme Diol +periodate

kr + diester kb

Values of the rate coefficients kb and kf have been determined for ethane-1,Zdiol andmethyl-substituted ethane-1,2-diols (Buistetal.”). The general effect of methyl substitution is to decrease k,. For meso- and (f)-butane-2,3-diols the values of kf are practically identical, although the equilibrium constant for forming the diester is much larger for the latter. This result suggests that formation of a monoester as an intermediate is rate determining, viz. References pp. 489-492

444

O X I D A T I O N O F O R G A N I C COMPOUNDS CH3CH-OH

I

CH,CH-OH

+

10;

I

CH3CH-OH

C H '~ C 0 -H -

'

fast

CH,CH--O

0- O 'H

0

0-

The steric effects responsible for the different equilibrium constants would be absent in the formation of the monoester. The kinetics of formation of the propane-1,2-diol diester at 25 "C have been studied over the range p H 1-13 by the stopped-flow technique (Buist and Buntoni3). The reaction is general base catalysed and shows many of the features found in the kinetics of the oxidation of pinacol (next section). The overall rate of oxidation of 2-methyl-butane-2,3-diolat pH 4.5 (acetate buffer) and 0 "C exhibits complex kinetics which can be analysed in terms of the diester being formed at a rate comparable with its rate of breakdown to the products (Buist and Bunton2'). The periodate species responsible for attacking 1,2-diols are probably lo;, and, in acid solution, H,I06 (Buist et aLZ6).Tracer studies show that the carbon-oxygen bonds in 1,Zdiols are not broken during the oxidation (Bunton and Shiner").

1.3.6 Kinetics of the oxidation of pinacol The reaction with pinacol is second order (first order with respect to each reactant)in therangepH0-11 (PriceetaZ.'9p28,Dukeand BulgrinZ9,Buist etaE.26*30). In the same pH range there is no evidence for the formation of an intermediate in appreciable concentration. The dependence of the second-order rate coefficient on pH is complex (Fig. 3) and shows maxima at pH 1 and pH 9.5. In the range pH 2-10 the reaction is general acid-base catalysed (Zuman et aL3', Buist et and for base catalysis the catalytic constants obey the Bronsted catalysis law with a slope of 0.69. Asnotedabove, the kinetics of pinacol oxidation resemble those

Fig. 3. Variation with pH of the second-order rate coefficient for the oxidation of pinacol by periodate at 25" and 0" C.

1

O X I D A T I O N S BY P E R I O D A T E

445

found for the formation of the diester from propane-1,2-diol and periodate. This ~ the mechanism of confirms the suggestion made earlier by Buist et ~ 1 . ' that pinacol oxidation is essentially the same as that for ethane-1,2-diol and lightly substituted diols, but that formation of the diester (or a monoester) is ratelimiting rather than the breakdown of the diester to the products. The general reaction scheme (3) includes a monoester, C1, as an intermediate in the formation of the diester C , , viz. D +Per

kt k z z C, e C , + Products k-1 k2

k-2

(3)

For ethane-1,2-diol the diester C , is in equilibrium with the reactants, and its decomposition to the reaction products is rate limiting and not subject to acidbase catalysis. When the concentrations employed are such that C2 is present in appreciable concentration, the mixed-order kinetics described in section 1.3.2 are observed. Second-order kinetics (for the overall reaction) can arise in three ways: (u) C 2 in equilibrium, but its concentration negligible, ( b ) formation of C , from C , rate limiting, and the latter in equilibrium with the reactants but present in very low concentration, and (c) formation of C , rate-limiting. For pinacol in the range pH 2-10 alternative (a) cannot be operative because general acid-base catalysis is observed. The most likely step to be subject to catalysis is the formation of Cz from C , , i.e. alternative (b), because this is a cyclisation and a base (B) could well facilitate reaction by removal of the C-OH proton, viz.

Furthermore, if alternative ( b ) is assumed the pH dependence of rate in the range pH 2-10 is explicable. The maxima at pH 1 and pH 9.5 are probably due to transitions to alternatives (a) or (c). With most catalysts in the pinacol reaction the normal linear dependence of rate coefficient on catalyst concentration is observed, but for the relatively strong base ammonia the dependence is markedly non-linear, and at high ammonia concentrations the rate coefficient approaches a limit. The non-linear dependence has been treated quantitatively by assuming a transition from alternative (b) to (c). Because the latter alternative involves a rate limiting step which is not catalysed, the limiting rate approached at high ammonia concentrations is explained (Buist et ~ 1 . ~ ' ) . In the range pH 9-11 the pinacol reaction deviates from strict second-order kinetics; the second-order rate coefficient decreases with increase of periodate References pp. 489-492

446

OXIDATION OF ORGANIC COMPOUNDS

concentration. This effect is more marked at 0" than at 25 "C, and it is due to the formation of the dimeric periodate species HzIzOIo4- (section 1.2) which is unreactive towards pinacol.

1.3.7 Kinetics of oxidation of other 1,2-diols (u) Methyl substituted cyclohexane-ly2-diols

Bunton and Carr3' found second-order kinetics for the oxidation of the cis1-methyl diol and the cis-lY2-dimethyldiol at pH 5.2 (acetate buffer). The lattLr diol is oxidised much more slowly, but ammonia is an effective catalyst (Bunton and Carr33). The observation of base catalysis shows that the reaction is of the pinacol type. The trans-lY2-dimethyldiol is almost inert to periodate, a result which is understandable on the basis of the steric effects discussed in section 1.3.4. (b) Cyclopentane-l,2-diols The oxidations of cis- and trans-cyclopentane-1 ,2-diols are second-order in the range pH 0-6 (Bulgrin and Dahlgren34) and exhibit a rate maximum in the region pH 1-2. These features resemble the kinetics of oxidation of pinacol (previous section) and the same conclusions regarding the mechanism can be drawn. The oxidations of the 1-methyl- and 1,2-dimethyl-cyclopentane-1,2-diols have also been studied by Bulgrin and D a h l g r e ~ and ~ ~ ~the , results again resemble those obtained for pinacol. The effect of methyl substitution on rate is shown in Table 2. Trans-l,2-dimethyl-cyclopentane-1,2-diolis inert. Qualitatively, the TABLE 2 S E C O N D - O R D E R R A T E C O E F F I C I E N T S F O R C Y C L O P E N T A N E - 1 , 2 - D I O L S AT I N UNBUFFERED SOLUTION

k2(Z.mo1e-'.sec- ')

pH 4-6 (0 "c)

cis-

trans-

cis-I-methyl

trans-I-methyl

cis-I,Ldimethyl

17.7

1.01

2.78

0.0112

0.0048

sequence of rate coefficients is in accordance with the steric factors operative in the diol-periodate cyclic esters: ( i ) greater ring strain for the trans-compounds, and (ii) steric compression between methyl groups and periodate oxygens (as in the model of the ester, Fig. 2). Evidently for the trans-1,2-dimethyl diol the combined effect of these two factors makes formation of the ester impossible. The increase in rate with increase of temperature for all the cyclopentane-1,2diols in Table 2 is small and corresponds to apparent activation energies in the range 1-5 kcal.mole-'. Each rate coefficient is probably the product of the rate coefficient for decomposition of the cyclic ester and one or more equilibrium constants. The latter could well decrease with increase of temperature, and bring about low overall activation energies.

1

O X I D A T I O N S BY PERIODATE

447

( c ) Some erythro-lthreo- epimeric pairs of the type RCH(OH)CH(OH)R'

Zuman et ~ 1 . studied ~' the oxidation of a number of such pairs of diols in the range pH 2-8 at 25 "C, and used polarography to follow the reaction. Dilute M ) of both reagents were used, and second-order kinetics solutions (ca. found in all cases with a dependence of rate on pH resembling that found for pinacol (section 1.3.6). Geneial acid-base catalysis was observed, and it is probable as with pinacol, that the formation of a diol-periodate ester is the rate limiting stage in the oxidation of all the compounds studied. Zuman et ~ 1consider . ~ that ~ the rate-limiting stage is a proton transfer step immediately prior to cyclisation, rather than cyclisation itself. The rate coefficients for threo- epimers arc greater than those for the corresponding erythro- epimers by factors varying from 1.2 to 4. Because the identity of the rate-limiting stage is uncertain, any interpretation of the rate differences is speculative, but it may be noted that the erythro- epimers are those for which steric compression in the cyclic diol-periodate ester must be greater. (d) Pyanosides

Honeyman and Shaw3' found second-order kinetics for the oxidation of a number of pyranosides, mostly at pH 4.1 (acetate buffer) and 25 "C. The pyranoside methyl-4,6-O-benzylidene-a-d-altroside, containing the two hydroxyl groups held rigidly in axial positions, is not oxidised by periodate (section 1.3.1). As with the cyclohexane-l,2-diols, those pyranosides containing cis-hydroxyl groups are generally oxidised more rapidly than those containing trans-hydroxyl groups, but without definite identification of the rate-limiting stage the interpretation of the rate differences remains uncertain. Honeyman and Shaw's work has bsen extended by Barlow and G ~ t h r i e ~ to~ ' include pyranosides not investigated by the former authors. (e) Miscellaneous Angyal and Young36 report second-order kinetics for the oxidation of the camphane-2,3-diols. The cis- isomers are oxidised much more rapidly than the trans- isomers, and a temperature of 80 "C had to be used for kinetic measurements on the latter. It seems likely that the rigidity of the camphane skeleton prevents the formation of a cyclic diol-periodate ester from the trans- isomers. Possibly the reaction at 80" is completely different in nature from the normal oxidation of 1,2-diols by periodate. The same workers37 report that cholestane-3P,6p,7atriol, in which the 6p and 7 a hydroxyl groups are axial-axial, is inert towards periodate. Senent-Perez et ~ 1 found . ~second-order ~ kinetics for the oxidation of tartaric acid over the range pH 1-12 with rate maxima at pH 4.5 and pH 7.5. Verma and Grover3' also found second-order kinetics, but with a single rate maximum References pp. 489-492

448

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

at pH 6.6. The discrepancy between the two groups of workers is probably due to the effects of buffer catalysis.

1.4

KINETICS OF O X I D A T I O N O F 1 , 2 - A M I N O A L C O H O L S

The stoichiometry of the oxidation of 1,2-arninoalcohols by periodate is RzC(OH)C(NHz)Ra+IO;

=

R2C-0

+ RaC-0 +NH, + 10;

Probably the initial products of the oxidation are a carbonyl compound and an imine, then the latter is rapidly hydrolysed, uiz. R,C(OH)C(NH,)R; +IO; = RzC-O

+R,C=NH +10; + H2O -1 H2O R;C-0 + NH,

The presence of an imine intermediate has not been demonstrated experimentally. If the amino group is secondary a similar oxidation occurs, e.g. CHzOH

I

CH2-N<

0 , = 2H2C=0 + CH3NH2+ 10; CH3+ 1 H

If the amino group is tertiary then any reaction that occurs is very slow (the formation of an imine is impossible in this case). Nearly all the rate studies of 1,2-arninoalcohol oxidations reveal second-order kinetics (first-order with respect to each reactant) but for phenyl-a-piperidylcarbinol (VI)

PI

OH

mixed order kinetics of the ethane-l72-diol type were observed by Kovar et al.40. This result is consistent with the formation of an appreciable concentration of an intermediate which is presumably a cyclic periodate ester. Several workers (Kovar et aL4', McCasland and Smith41, Dahlgren and H o d ~ o nhave ~ ~ )shown that the second-order rate coefficients for the oxidation of 1,2-aminoalcohols increase with increase of pH in the range pH 3-9. The rate coefficient is proportional

1

449

O X I D A T I O N S BY P E R I O D A T E

to the concentration of free amine, and it is usually assumed that the latter is the reactive species, and that the protonated amine is not oxidised. The rapidity of the oxidation at high pH values has confined most kinetic studies to the region pH > 7, or to low concentrations ( M ) of the reactants. Another general feature of the oxidation of 1,2-aminoalcoholsby periodate is the small temperature coefficient, e.g. McCasland and Smith4' found that the rates of oxidation of the 2-amino-cyclohexanols and 2-amino-cyclopentanolswere only reduced by a factor of two when the temperature was changed from 25" to 0°C. Dahlgren and Hodson4' calculated rate coefficients assuming the following rate equation: N

(Per = total periodate, and A = free amine). They found that k2 decreases with increase of temperature (0-25 "C). Dahlgren et ~ 1 .44~ point ~ . out that if H,IO, is assumed to be the reactive species rather than IO;, then k2 has a positive temperature coefficient. However, all interpretations of the kinetic data are uncertain as yet, because the rate-limiting stage has not been identified with any certainty. Dahlgren and Rand44 observed general base catalysis of the oxidations of 2-amino-ethanol and of N-methyl-2-amino-ethanol,so it is possible that in these cases ring closure to a cyclic periodate ester is the rate-limiting step (cf. the oxidation of pinacol, section 1.3.6). Kovar et a1.4' determined the second-order rate coefficients for a number of cisltrans and erythrolthreo isomeric pairs of 1,2-aminoalcohols, including some with a methyl or benzyl group on nitrogen. In most cases the rate of oxidation of the trans- or the threo- isomer is faster. Methyl substitution on nitrogen enhances the rate difference between members of a pair. However, benzyl substitution on nitrogen reduces the rate difference, and, in some cases, reverses the order of reactivity, making the rate of oxidation of cis- or erythro- isomers greater by a factor between 2 and 40. Kovar et al. ascribe this effect to hydrogen bonding TABLE 3 SECOND-ORDER

(l.mole-'.sec-')

F OR T H E OXIDATION OF SOME

1,2 A M I N O AL COHOLS A N D 25 OC

A T pH6.8 ( P H O S P H A T E BUFFER)

R A T E CO E F F I CI E NT S

cisltrans A N D e r y t h r o l m e s o

PAIRS OF

2-amino-cyclohexanol N-methyl-2-amino-cyclohexanol N-benzyl-2-amino-cyclohexanol 1,2-dipheny1-2-amino-ethanol N-methyl-l,2-diphenyl-2-aminoethanol N-benzyl-3-aminopentane-2,5-diol References pp. 489-492

cis

trans

1.6

5.2 18 0.04

0.77

1.6

crytho

3.7 0.88 14

threo

50

16

6.0

450

O X I D A T I O N OF O R G A N I C C O M P O U N D S

between the hydroxyl hydrogen and the aromatic ring. The hydrogen bonding is expected to be stronger for the cis- or erythro- isomers, and is assumed to hold the molecule in a more favourable conformation for reaction with periodate. Some of the rate coefficients obtained by Kovar et aL4' are listed in Table 3. In contrast to the results of Kovar et al., McCasland and Smith4' found that cis-2-amino-cyclopentanol is oxidised more rapidly than the trans- isomer. Barlow et ~ 1determined . ~ ~the second-order rate coefficients at pH 4.0 (acetate buffer) and 25 "C for a series of amino sugars derived from the eight isomeric methyl-4,60-benzylidene-cr-D-glycosides(VII)

ph cdo/2-+ om o\ x2

X'. X',Y: and Y 2 = H /OH /NH,

Y2

mr Each rate coefficient was calculated for the reaction of the free amino group, thus the different basicities of the amino sugars were allowed for. Generally those amino sugars having a cis arrangement of amino and hydroxyl groups are oxidised more rapidly. However, the cis compound, methyl4,6-O-benzylidene-aD-3-amino-alloside, is oxidised about four times more slowly than the compounds having an eq-eq trans arrangement of NH2 and OH. The ax-ax trans- compounds, the amino-altrosides, are not inert, but are oxidised very slowly in a phosphate buffer at pH 6.9 (cf. the inertness of methyl-4,6-O-benzylidene-cl-~-altrosides, section 1.3.7. d). It is evident that the order of reactivity of cis and trans isomers is variable. Their reactivities may well be related to the ease of formation of a cyclic periodate ester, but again any interpretation is limited by uncertainty concerning the rate limiting stage. Barlow and GuthrielS6* 5 7 later reported their work in more detail, and included kinetic data for the pH range 4-7, at 0" as well as at 25°C. They suggest that the amino-altrosides are oxidised by an unknown mechanism, different from the mechanism of oxidation of the other six isomers.

1.5

KINETICS OF OXIDATION OF

1,2-DIKETONES

AND

1,2-HYDROXYKE-

TONES

1.5.I Oxidation of I,2-diketones

The few kinetic studies of the oxidation of 1,2-diketones indicate that the mechanism differs from that found for the oxidation of 1,2-diols, and that, contrary to earlier suggestions, the reaction does not proceed uia a hydrate of the diketone. Shiner and W a ~ m u t hstudied ~~ the oxidation of butane-2,3-dione,

1

45 1

OXIDATIONS BY PERIODATE

TABLE 4 RATE

COEFFICIENTS

FOR

THE

OXIDATION

OF

SOME

1,2

DIKETONES

BY

VARIOUS

P E R I O D A T E SPECIES

Rate coefficients* (l.mole-'.sec- ') Temp.

("C) Butane-2,3-dione** 2,5-Dimethylhexane-3,4-dione** Benzil** Camphorquinone** Glyoxalt Pyruvaldehydet Butane-2,3-dionet

*

k-

kZ-

0.74

13.3

350 0.61 3.1 0.52 5000 3000 639

1.5 11.8 1.46

714 330 15.7

k314 26 18

-

Rate coefficients defined by

where [Per]

**

22 22 22 22 26 26 26

k"

= total

periodote and K

=

diketone.

Shiner and W a ~ r n u t h ~ ~ .

t Dahlgren and Reed43.

2,5-dimethyl-hexane-3,4-dione,benzil, and camphorquinone, over the range pH 1-12 (buffered solutions) at 0" and 22 "C. With each compound the kinetics are strictly second order (first order with respect to each reactant) and above pH 7 the rate increases markedly with increase of pH. The reaction is slightly buffercatalysed, and the catalysis is probably general base only. The dependence of rate on pH is accounted for fairly well by assuming second-order reactions with each of the periodate species H5106, H,IO;, H,IO;-, and H2102-. Table 4 lists the rate coefficients for reaction of these species, and includes similar data obtained by Dahlgren and Reed43 for the oxidation of glyoxal, pyruvaldehyde and butane2,3-dione. The marked drop in rate coefficients in going from pyruvaldehyde to butane-2,3-dione can be explained by a steric effect (Dahlgren and Reed43). From a study of the oxidation of the oxidation of " 0 labelled butane-2,3dione, Bunton and Shiner27 concluded that the reaction does not proceed via the dihydrate of the dione (at least in the alkaline conditions which had to be employed for the " 0 experiment). Reaction of the monohydrate could not be ruled out, but it appears more likely that reaction occurs via nucleophilic attack of periodate on both carbonyl groups. Moreover, the rate coefficients in Table 4 follow the expected order of nucleophilicity. Note that if 10; is assumed to be reactive instead of H,IO;, the order of reactivity for H510, and 10; is wrong for nucleophilic substitution (10, must be a weak nucleophile). Also, a very small activation energy is found for reaction assuming IO,, but for H410; the activation energy is about 10 kcal.mole-', the same as for reaction with H5106 References p p . 489-492

452

OXIDATION OF O R G A N I C COMPOUNDS

(Dahlgren and Reed43). Shiner and W a ~ m u t hsuggest ~ ~ that the oxidation of 1,Zdiketones proceeds via a cyclic periodate ester, similar to that postulated for the oxidation of 1,2-diols, but formed by nucleophilic attack of periodate on carbon, uiz. R

+

HJO;

R-C-C-R'

l

o

I

0-

II II

0

0

RCOOH

+

R'COOH

+ 10; +

H,O

The actual presence of a cyclic periodate ester or other intermediate has not been demonstrated experimentally. Dahlgren and Reed43 suggest that the general base catalysis could be due either to an interaction of the catalyst with periodate in a prior equilibrium, or to proton abstraction prior to formation of the cyclic ester.

1.5.2 Oxidation of 1,bhydroxyketones Only one kinetic study has been reported. The oxidation of 3-hydroxy-3methyl-butane-2-one is second-order with a rate maximum at pH 8 (Bunton and Shiner"). The same authors carried out an l 8 0 experiment at pH 8 which shows that the C-OH bond is not broken, as in the oxidation of 1,2-diols. The mode of attack on the carbonyl group is probably nucleophilic at the carbon atom, as in the oxidation of 1,2-diketones. 1.6

K I N E T I C S O F O X I D A T I O N O F QUINOL A N D CATECHOL,

THEIR MONOMETHYL ETHERS, A N D ESTERS OF Q U I N O L

Phenol itself is only oxidised slowly by periodate (Adler and Magnusson4') but many substituted phenols are oxidised rapidly. Sklarz4 has reviewed the extensive work done on the stoichiometry of the oxidations of the latter. Few kinetic studies have been Garried out. Kaiser and Weidman48found second-order kinetics (first with respect to each reactant) for the oxidation of quinol and its monomethyl ether at 15" and 25 "C,i.e.

1

453

O X I D A T I O N S BY P E R I O D A T E

At pH 1.0 the second-order rate coefficients are expressed by k,(quinol) k,(ether)

=

=

1.0 x 10" exp (- 11,20O/RT)l.mole-'.sec-'

2.5 x lo9 exp (- l1,000/RT) l.mole-'.sec-'

Increase of pH in the range 0-4 decreases the rate. The pH dependence can be accounted for quantitatively by assuming different reactivities for H5106 and a periodate monoanion. Kaiser and Weidman49. later studied the oxidation of catechol using a stopped-flow apparatus, and showed that the reaction proceeded via an intermediate. The formation of the latter was second order, and at pH 1.0 the second-order rate coefficient is expressed by (from measurements at 15" and 25 " C ) k,

=

6.4 x lo9 exp (- lO,SOO/RT) 1.mole-'.sec-'

The activation parameters are similar to those for the overall reactions of quinol and its monomethyl ether (above). The pH dependence of k2 shows a gradual increase in the range pH 1-4, then a more rapid increase to a maximum at pH 8.5, followed by a rapid decrease. Kaiser and Weidman accounted for this quantitatively by assuming reactions of catechol with H5106 and a periodate monoanion, and either reaction of catechol with H31062-, or reaction of the catechol monoanion with a periodate monoanion. The latter possibilities could not be distinguished. The decomposition of the intermediate in the catechol oxidation is first-order and is subject to catalysis by hydrogen ions, viz.

At 25 "C,k, = 0.24 sec-', kH = 26 1.mole-'set-', and koH = 2.6 x loB1.mole-'. sec-'. The latter coefficient is about 1/40 of the rate coefficient for a diffusion controlled reaction in water. Kaiser and Weidman' suggest that the intermediate is similar to the cyclic periodate ester postulated for the 1,2-diol oxidations, viz.

+ OH

10,-

-

+

10;

The apparent hydrogen and hydroxide ion catalysis can be accounted for by asReferences pp. 489-492

454

O X I D A T I O N O F O R G A N I C C OMPOU N D S

suming different rates of breakdown for the undissociated ester and its anions (cf. the situation for the cyclic ester formed from 1,2-diols; the monoanion is the only species whose decomposition has been detected, section 1.3.2). However, Kaiser and Weidmans O presented some evidence which suggests that the intermediate may be of a different type, e.g. a charge-transfer complex between 10; and o-benzoquinone. In the oxidation of guaiacol (the monomethyl ether of catechol) no intermediate can be detected (Kaiser and Weidman4'). When quinol or catechol are oxidised by periodate in H,"0, the benzoquinones produced are not labelled (Adler et aL5') showing that the carbon-oxygen bonds are not broken. However, the oxidations of the monomethyl ethers of quinol and catechol in H2"0 produce 50 % labelled benzoquinones and unlabelled methanol. Probably the "0 enters as shown in reactions (a) or (b)

H :,

+

CH,OH

Esters of quinol are oxidised by periodate much more slowly than quinol itself. Wiedman et al.' 5 8 studied the oxidation of p-hydroxyphenyl sulphate (quinol monosulphate) in aqueous solution at pH 1 and 25 "C.The rate is about 2000-fold slower than the oxidation of quinol under the same conditions. Because the rate depends on the periodate concentration, Wiedman et al. concluded that oxidative cleavage of the ester occurs, rather than hydrolysis followed by oxidation of quinol. Oxidation of the ester in H2"0 showed that less than 10 % of S - 0 bond cleavage occurs; therefore most of the sulphate is released as HSO, rather than SO,. However, in methanol as solvent monomethyl suIphate is formed in high yield, and under these conditions it is probable that the sulphate is released as SO,. Bunton and Hellyerl 5 9 studied the oxidation of p-hydroxyphenyl acetate and benzoate, 2,5-di-tert.-butyl-p-quinol, and 3-tert.-butyl-4-hydroxyphenylacetate in aqueous solutions in the range pH 0.5-3.0 at 25 "C. The oxidations are secondorder (first with respect to each reactant) and in the case of the esters the products are the p-quinone, iodate, and acetic or benzoic acids. The rates of oxidation of p-hydroxyphenyl acetate and benzoate are very similar. The steric effect of the tert.-buty1 group, substituted into p-quinol and p-hydroxyphenyl acetate, is to reduce the rate by almost the same factor in both cases. Bunton and Hellyer concluded that the mechanisms of oxidation of quinol and its esters by periodate are the same, and that the rate-determining step is the electrophilic attack of either

1

O X I D A T I O N S BY P E R I O D A T E

455

H,I06 or 10; on the hydroxyl group to form a periodate ester whose rate of decomposition is fast. The same workers'60 also studied the oxidation of p-hydroxyphenyl phosphate by periodate (aqueous solutions, pH 0-10,25"C). The kinetics are second-order, and the rate maxima are at pH's 0.7 and 7.5. Bunton and Hellyer accounted for the pH dependence by assuming (a) that both the ester and its monoanion react with HSI06, and (b) that the ester monoanion and dianion react with a pzriodate monoanion. The decrease in rate beyond pH 7.5 shows that reactions involving the phenoxide ion or the periodate dianion are unimportant. 1.7

OXIDATION OF ACTIVE METHYLENE GROUPS

Huebner et al. 5 2 studied the stoichiometry and approximate rates of oxidation of a number of compounds containing an active methylene group. They found that not all such compounds are oxidised by periodate, and that in general one of the activating groups must be -CHO or -COOH for oxidation to occur. Thus diethyl malonate, ethyl acetoacetate, and cyanoacetic acid are not oxidised. Acetylacetone and other acyclic lY3-diketonesare oxidised very slowly, but cyclic lY3-diketonesare readily oxidised (Wolfrom and B ~ b b i t t ~The ~ ) . first step in the oxidation of a compound containing an active methylene group is hydroxylation, then this is followed by further oxidation, e.g. malonic acid and 1,3-~yclohexanedionereact as follows CH2(COOH )2

I0,CHOHKOOH),

Io4- CHO(CO0H) -

4- COP

110,HCOOH

+

CO2

The only compound containing an active methylene group whose periodate oxidation has been studied kinetically is 2-methyl-l,3-cyclohexanedione(Wolfrom and B ~ b b i t t ~At ~ )22.5 . "Cand pH 6.2 the kinetics are first-order with respect to each reactant. The same workers showed qualitatively that for the oxidation of 5,5-dimethyl-l,3-cyclohexanedione,the rate reaches a maximum between pH 5 and 6 (phosphate buffers were used). References p p . 489-492

456

O X I D A T I O N OF O R G A N I C C O M P O U N D S

The pH dependence may be due to the reactive periodate species being IO;, but the mechanism of hydroxylation is uncertain. The exceptions noted above show that enolisation cannot be the sole factor determining whether or not hydroxylation occurs, furthermore some weakly enolised compounds (e.g. malonic acid) are readily oxidised. Bose et al. 54 suggested a cyclic mechanism, but such a mechanism cannot be extended readily to malonic acid, or, for steric reasons, to 1,3-cyclohexanedione (Sklarz4). Bunton5 has suggested that hydroxylation may occur by 10; acting as an electrophilic oxidant transferring oxygen to the substrate, uiz. >CH2

+ o=xo;

’-

‘C“0H /

+

10;

Recently, Yadawa and Krishna161 studied the oxidation of malonic acid by periodate at 30” C in the range pH 2-8. The kinetics are second-order, and the rate reaches a maximum at pH 6. Yadawa and Krishna concluded that the malonate dianion is reactive, but their data are insufficient to allow details of the mechanism to be inferred. 1.8

P E R I O D A T E O X I D A T I O N S I N MIXED S O L V E N T S

Periodate oxidations are much slower in non-aqueous solvents or in mixed aqueous-organic solvents. Very few kinetic studies have been made of the solvent effect. Taylor et al.” found that the rate of oxidation of ethane-1,Zdiol by periodic acid is reduced by a factor of approximately 12 by the addition of 38 % of ethanol to water. Guthrie’ studied the oxidation of trans-cyclohexane-l,2diol in dimethylformamide/water mixtures, and found a sharp reduction in rate when the DMF content was increased beyond 20 %. 2. Kinetics of oxidations by peroxodisulphate 2.1

I N T R O D U C T I O N (see also

Chapter 4)

2.1.1 General properties and spontaneous decomposition of peroxodisulphate Peroxodisulphuric acid, H2S20s, is a strong acid whose second pK is below zero (Kolthoff and Miller”). Under the conditions normally employed in peroxodisulphate oxidations (aqueous solution, pH > 1) the ion SzOs2-is the dominant species. The ion is a powerful two-electron oxidising agent with a redox potential of -2.01 V. In the majority of its reactions the primary step is the formation of sulphate radical-ions, either by spontaneous fission of the peroxide bond, or by attack on a substrate X, i.e.

2

457

OXIDATIONS B Y PEROXODISULPHATE

or s208’- +X -+ SO;

+ so4’- +x+

(2)

The highly reactive sulphate radical-ion may attack the solvent or a substrate present in the solution. The spontaneous decomposition of peroxodisulphate in aqueous solution has been widely studied; it is, in effect, the oxidation of the solvent by peroxodisulphate. It will be considered here because the steps involved are of importance in the mechanisms of oxidation of organic compounds by peroxodisulphate. The decomposition has the stoichiometry S20S2-+H20 = 2 H + + %0 2 + 2 SO4’’

(3)

Early work established that the kinetics of decomposition are first order with respect to peroxodisulphate, and in a detailed study Kolthoff and Miller” showed that the rate equation includes an acid-catalysis term, viz.

The values of ko obtained by Kolthoff and Miller for the reaction in 0.1 M sodium hydroxide in the range 50-90 “Care given in Table 5 . TABLE 5 R A T E COEFFICIENTS F O R T H E DE CO M P O S I TION OF

S20sz-

IN

0.1 M NaOH

Activation energy = 33.5 kcalmole-’ Temp. (“C)

50

lo5 ko(sec-’)

60 0.50

0.10

70

2.3

80

9.2

90

35

In their discussions of the decomposition, Tsao and Wilmarth6’, and Wilmarth and Haim”, propose the following mechanism for the uncatalysed reaction (based on earlier suggestions by Kolthoff and Millers9)

2 OH -+ H 2 0 2 H202 References p p . 4 8 9 4 9 2

=

H20+$02

(7)

458

O X I D A T I O N OF O R G A N I C C O M P O U N D S

S20i-+H202 = 2 H + + 2 S0,Z-+02

(9)

(details of reactions (8) and (9) are not shown). The first step is assumed to be ratedetermining, in accordance with the first-order kinetics. Tracer studies using 3 5 S labelled sulphate show that exchange of sulphur between sulphate and peroxodisulphate is negligible except at relatively high concentrations of hydrogen and sulphate ions (Tsao and Wilmarth60, and references cited therein). The observed exchange can be explained by reaction (6), whose reverse step is only important when the product [H+][S042-] is sufficiently great, and by reaction (10) which is only important when the concentration of sulphate radical-ions is relatively high. Results of H2180 experiments (Kolthoff and Miller59, Tsao and Wilrnarth6') show that the evolved oxygen originates from the solvent, in accordance with the above mechanism. However, the mechanism is not entirely satisfactory. It does not explain the variation of ko with pH in unbuffered solutions (Breuer and Jenkins6'); also Fronaeus and Ostman62'63 have put forward evidence favouring reaction (11) as the first step, rather than reaction ( 5 ) , viz.

Another alternative ( L e ~ i t t involving ~~) rapid, reversible formation of sulphur tetroxide and sulphate radical-ions, is ruled out by the 5S exchange experiments. shows that when peroxodisulphate is allowed Recent work by Crematy'62' to decompose in the presence of a free-radical scavenger, the initial rate of production of acid is zero. This observation is incompatible with reaction (ll), and Crematy concluded that the first step is reaction (5). In strongly acid solutions ([H'] > 0.5 M ) the major products of the decomposition are peroxomonosulphuric acid and hydrogen peroxide (Kolthoff and viz. H2S208+H20 = H2SOS+H2S04

+

H2S05 H 2 0 = H 2 0 2+ H2S04

(12) (13)

The mechanism of decomposition under these conditions is uncertain, but probably HSO; is the initial product (Wilmarth and Haim57). Many of the reactions of peroxodisulphate are catalysed by silver ions, and although the reactions then involve higher oxidation states of silver as oxidants, it is convenient to consider them along with the uncatalysed oxidations. Silver ions also catalyse the decomposition of peroxodisulphate in a second-order reaction, viz.

2

OXIDATIONS BY PEROXODISULPHATE

459

The mechanism is uncertain. Bawn and Margerison6 proposed the following S20i-+Ag' Ag2++OH-

-+ -+

Ag2++SO,

+ SO:-

Ag+ +OH

(15) (16)

The sulphate radical-ions and the hydroxyl radicals undergo further reactions as in the mechanism for the uncatalysed decomposition. The formation of siIver(II1) ions has also been proposed (Yost66,Morgan67).

2.1.2 Types of peroxodisulphate oxidation mechanisms Oxidations of both inorganic and organic substrates by peroxodisulphate have been reviewed by Wilmarth and Haim57and by House5'. Peroxodisulphate oxidises a variety of organic compounds, including alcohols, aldehydes, ketones, hydroxy-acids, phenols, and amines. The majority of the oxidations proceed via free-radical chain mechanisms, but a few are ionic in character, notably the oxidations of phenols and aromatic amines, and involve nucleophilic displacement on the peroxide oxygen. The chain mechanisms are characterised by the following: the reaction proceeds at a faster rate than the spontaneous decomposition of peroxodisulphate, fractional orders with respect to peroxodisulphate and the substrate are often encountered, and the reactions are often susceptible to catalysis or inhibition by impurities (particularly metal ions and dissolved oxygen). Reproducible kinetics are often difficult to obtain, and the interpretation of much of the available data is uncertain due to the use of solutions containing unknown amounts of oxygen or metal ions. The possible types of chain mechanisms for peroxodisulphate oxidation have been classified by Wilmarth and Haim57 according to the dominant initiation and termination steps, and the relative importance of sulphate radical-ions and hydroxyl radicals in the propagation steps. Some of the rate equations corresponding to the different types of mechanisms are the same, so the observation of a particular rate equation does not always permit a unique mechanism to be inferred. In certain cases the nature of the chain initiation step can be deduced from the effect of a free-radical scavenger on the reaction rate. Thus in the oxidation of 2-propanol, the addition of ally1 acetate reduces the rate to that observed for the spontaneous decomposition of peroxodisulphate, indicating that the chain initiation step is the same as the rate-determining step of the spontaneous decomposition, viz. the fission of peroxodisulphate into sulphate radical-ions. As noted above, many peroxodisulphate oxidations are catalysed by silver References pp. 489-492

460

O X I D A T I O N OF O R G A N I C C O M P O U N D S

ions. With the majority of inorganic substrates, such oxidations are first order with respect to both peroxodisulphate and silver ion. The rates are independent of the concentration and nature of the substrate, showing that the oxidations have a common rate-determining step involving peroxodisulphate and silver ions only. This step is usually assumed to be reaction (15). With organic substrates the rates are usually faster and vary considerably from one substrate to another. Fractional orders with respect to peroxodisulphate and silver ions are common, but the rates are usually independent of the substrate concentration. Undoubtedly chain mechanisms are operative in the majority of cases, but with few exceptions mechanistic details are uncertain. In addition to silver@) ions and sulphate radical ions, silver(II1) ions and hydroxyl radicals may participate in the oxidation, and further complications are introduced if complexing occurs between the substrate and silver ions. In the following sections the peroxodisulphate oxidations of organic compounds are classified according to the nature of the substrate. Unless otherwise stated the solvent is water. 2.2

O X I D A T I O N S OF ALCOHOLS A N D DIOLS

The general stoichiometry of the oxidation of alcohols is

R,R2CHOH + SzO:- = RIRzC-O+ 2 SO:- $2 H+

(17)

Ri and R2 can be alkyl groups or hydrogen. The oxidation of one tertiary alcohol, tert.-butyl alcohol, has been studied; the stoichiometry is uncertain, but acetone and formaldehyde are the main products (section 2.2.5). 2.2.1 Oxidation of 2-propanol

It is convenient to discuss this oxidation first because it has been investigated more thoroughly than the oxidations of other alcohols. Levitt and Malinowski6**6 9 showed that the reaction is first-order with respect to peroxodisulphate, and found that the rate is independent of the concentration of 2-propanol at high concentrations of the latter. They proposed an ionic mechanism involving the reversible formation of an ester, uiz.

(CH,),CHOH+ SzOi'

+

+ (CH,),CH-O-0-SO; I 1

+SO:-

H

(CH,),C-O+2 H++SOZ-

2

46 1

OXIDATIONS B Y PEROXODISULPHATE

Later work showed this mechanism to be incorrect. Wiberg” showed that 3 5 S 0 2 - when present in the reaction mixture does not give labelled peroxodi-

sulphate, as required by the reversible first step of Levitt and Malinowski’s mechanism. Furthermore, allyl acetate inhibits the reaction and reduces the rate of consumption of peroxodisulphate to that observed in the absence of 2-propanol. Wiberg proposed a chain mechanism involving sulphate and hydroxyl radicals. In a thorough study of the reaction, Ball et al.” showed that all previous studies were complicated by the catalytic effects of trace amounts of metal ions (most likely cupric ions) and inhibition by dissolved oxygen from the atmosphere. In the absence of oxygen there is no catalysis by cupric ions, and the rate equation is

- dCSz0a2-1 dt

= k[S2082-][(CH3)2CHOH]~ TABLE 6

R A T E C OE FFI C I E NT S FOR T H E O X I D A T I O N OF 2 - P R O P A N O L

BY

PEROXODISULPHATE

I N T H E AB S E NCE OF O X Y O E N

Activation energy = 21.0 kcal.mole-’ Temp. (“C)

lo2 k(lhnole-ksec-’)

45 0.50

50

0.77

55 1.27

60 2.42

65

3.57

70 4.75

Table 6 lists values of the rate coefficient k obtained by Ball et al. They proposed the following chain mechanism

szo;- 1:2 so,

(20)

2(CH,)$OH+HSO; ( C H 3 ) , t 0 H + S Z 0 ~ -2(CH3)zC-O+HSO; (CH,),COH+ SO; 2 (CH,),C-O + HSO;

(21)

SO; +(CH3)&HOH

+SO,

(22) (23)

The stationary-state approximation applied to this mechanism gives a rate equation agreeing with equation (19), viz.

- dCSzdtOa2-’ = (k,kz k3/k4)’[S2Oa2-] [(CH,),CHOH]*

(24)

As noted in section 2.1.2, the inhibition by allyl acetate is ascribed to the removal of sulphate radical-ions formed in the initiation step. The inhibited rate is 1800 times smaller than the rate in the absence of allyl acetate, indicating a chain References pp. 489-492

462

O X I D A T I O N OF O R G A N I C COMPOUNDS

length in excess of lo3. From equation (24) the observed activation energy, EA, is related to the activation energies of the elementary reactions by the equation

El is approximately 34 kcal.mole-I (section 2.1.1), and the other activation energies must be small because the chain length is great. Therefore the observed activation energy (21 kcal) is in qualitative agreement with equation (25). Dogliotti and H a y ~ n ’generated ~ sulphate radical-ions at room temperature by the flash photolysis of pxoxodisulphate solutions, and measured the rate at which they attack 2-propanol, i.e. ths rate of reaction (21). They found k2 = 8.5 & 3.0 x lo7 1.mole.sec-’ at pH 4.4. In the presence of oxygen and absence of metal ions (Ball et al. added EDTA to the reaction mixture to reduce trace amounts of free metal ions to insignificance) an induction period is observed whose length depends on the oxygen partial pressure. A chain mechanism involving reactions (26)-(28) in addition to reactions (20) and (21) is in qualitative agreement with the observed induction periods. (CH3),COH

+ O2 2.00(CH3),COH

(26)

2

*OO(CH3)2COH+(CH3)2CHOH HOO(CH3),COH+(CH3),~OH(27) 2 0O(CH3),COH

1:HOO(CH,),COH+(CH,),C-O+

0,

(28)

Ball et aL7’ also studied the cupric ion-catalysed reaction (in presence of oxygen) and found that the rate is independent of the 2-propanol concentration; the expression being

The catalysis has an upper limit at approximately [Cu”] = M. The mechanism is uncertain, but Ball et al. propose the following steps in addition to reactions (26)-(28)

+

2Cu3*+HOO(CH,),COH

(30)

4 Cu++(CH3),C=O+2 H +

(31)

H + 00(CH3)2COH+C~2’ Cu3++(CH3)CHOH

cu++szo82-2 cu3++2 so:-

i32)

2

OXIDATIONS BY PEROXODISULPHATE

2 cu+

2 cuo+cu2+

c u 0 + s 2 o ~ -Y c u 2 + + 2 so,"-

463 (33) (34)

The stationary-state approximation applied to the mechanism gives the rate equation - dCS2082-1 =

dt

(k,/k,,)9k,o[S20s2-])

(35)

which is in agreement with equation (29). In a study of the silver ion-catalysed oxidation of 2-propanol in 50% acetic acid, Venkatasubramanian and S a b e ~ a nfound ~ ~ first-order kinetics with respect to peroxodisulphate and zero-order with respect to the alcohol. They propose a chain mechanism involving either silver(I1) or silver(II1) ions.

2.2.2 Oxidation of methanol and ethanol

Bartlett and C ~ t m a found n ~ ~ that the kinetics of oxidation of methanol obeyed the rate equation

- d[S2 OS2 dt

=

k [ S 2 02 -]*[CH3 OH]*

The reaction is faster than the spontaneous decomposition of peroxodisulphate, so Bartlett and Cotman proposed a chain mechanism involving sulphate radicalions and the CH20H radical. Kolthoff et showed that ally1 acetate inhibits the reaction, and reduces the rate to that observed in the absence of methanol. They pointed out that if the inhibition is explained on the basis of Bartlett and Cotman's mechanism, the predicted rate equation does not include the methanol concentration. This difficulty was resolved by Edwards et ~ l .who ~ ~showed , that in the absence of oxygen the reaction is zero-order with respect to methanol. They proposed the following mechanism (similar to that originally proposed by Bartlett and Cotman)

S04-+CH30H

2 CH20H+HSO4-

(38)

2

(39)

~ H , 0 H i - S 2 0 ~ - HCHO+HSO;+S04References p p . 489-492

464

O X I D A T I O N OF O R G A N I C COMPOUNDS

2 cHzOH

2products

(40)

For long chains (the observed chain length is 100 at 70 "C and initial [SzOi-] = M ) the stationary state approximation gives

Note that a different termination step from that in the 2-propanol oxidation is responsible for the rate being independent of the substrate concentration. The absence of kz in equation (41) predicts the same rate for CD30H. The observed isotope effect (kH/kD) is 1.3 which can reasonably be ascribed to a secondary effect on k3 or an inverse effect on k4 (Edwards et ~ 1 . ~ ~ ) . In the oxidation of ethanol, the order with respect to peroxodisulphate is again three-halves and the mechanism is considered to be the same as that proposed for methanol, but the oxidation is complicated by inhibition due to the product, acetaldehyde (Edwards et ~ 1 . ~If~ sufficient ) . acetaldehyde is added to the reaction mixture to keep its concentration essentially constant, then the observed rate is independent of the ethanol concentration. Otherwise, the rate increases with increase of initial ethanol concentration at a given peroxodisulphate concentration. Evidently acetaldehyde competes with the ethanol for the oxidising radicals, and Edwards et al. suggest the following scheme CH3CH0+ SO4-

-+

C H 3 c 0+ HS04-

C H 3 c O + S z O ~ - + H z 0-+ CH,COOH+HSO,+SO, C H 3 c 0+ CH3cHOH 3 products

(42) (43)

(4)

The stationary-state approximation leads to a rate equation in agreement with the observed expression. Dogliotti and H a y ~ n 'have ~ determined the rate of attack of sulphate radicalions at room temperature on methanol, i.e. reaction (38), and found k2 = 2.5k0.4 x lo7 l.mole-'.sec-' at pH 4.8. The corresponding rate coefficient for attack on ethanol is 7.7k2.2 x lo7 l.mole-l.sec-'. 2.2.3 Oxidation of 2-butanol

Subbaraman and S a n t a ~ p afound ~ ~ that with de-aerated solutions the rate equation for the oxidation of 2-butanol is the same as that found by Ball et aL7' for the oxidation of 2-propanol in the absence of oxygen, i.e. equation (19).

2

OXIDATIONS B Y PEROXODISULPHATE

465

Measurements in the temperature range 55-70 "C show that the rate coefficient is expressed by

k

=

6.64 x l o i 3exp (-25,20O/RT) 13.mole-3.sec-'

The mechanism is considered to be the same as that proposed by Ball et al., i.e. reactions (20)-(23). For the silver ion-catalysed oxidation, Subbaraman and S a n t a ~ p found a ~ ~ the rate equation to be

- d[S20i-1 dt

=

k,,[SzOi-][2-butano1]3[Ag+]f

(45)

However, Venkatasubramanian and S a b e ~ a nreport ~ ~ that the rate is zero order with respect to the substrate, as in most silver ion-catalysed oxidations by peroxodisulphate.

2.2.4 Oxidation of cycIohexano1 The uncatalysed oxidation in de-aerated solutions is first order with respect to peroxodisulphate, and zero order with respect to the substrate (Subbaraman and Santappa7'). Measurements in the temperature range 55-70 "C show that the first-order rate coefficient is expressed by

k = 1.47 x l O I 4 exp (- 27,6OO/RT) sec-' The proposed mechanism, reactions (46)-(50)

s20,2-+ 2 so, SO;+H20

OH + RCHOH RCOH+S,OiRCOH+SO,

(46)

+

OH+HSO;

(47)

-+

RCOH + H,O

(48)

RC=O+HSO, + S O,

(49)

-+

+ RC-O+HSO;

(50)

(R = -(CH2)5-) involves hydroxyl radicals as well as sulphate radical-ions. According to the classification given by Wilmarth and Haim57, this mechanism is the only one giving the observed rate equation. References p p . 489-492

466

O X I D A T I O N O F O R G A N I C COMPOUNDS

The silver ion-catalysed oxidation of cyclohexanol obeys the rate equation (Subbaraman and Santappa7')

A mechanism consistent with this rate equation is obtained by adding reactions (15) and (52) to the reactions proposed for the uncatalysed oxidation. (15)

Ag++SzOi- + Ag2++SO;+SO;-

+

Ag2+ RCHOH -P Ag'

+ RCOH + H+

(52)

2.2.5 Oxidations of other alcohols

Subbaraman and Santappa7' found that the uncatalysed oxidation of tert-butyl alcohol is very slow, but the silver ion-catalysed oxidation proceeds at a readily measurable rate which is independent of the alcohol concentration, viz.

(53) The stoichiometry shows some variations with the peroxodisulphate and alcohol concentrations. It is given approximately by 2 SzOz-+(CH3)3COH+H,O = (CH3)zC-O+H,C-O+4

HSO;

(54)

Subbaraman and Santappa suggest a chain mechanism involving hydroxyl as well as sulphate radicals. Venkatasubramanian and S a b e ~ a nstudied ~ ~ the silver ion-catalysed oxidations of 2-pentanol and 2-octanol in 50 % acetic acid. In both cases they found firstorder kinetics with respect to peroxodisulphate, and zero-order with respect to the substrate. They propose a chain mechanism involving silver(I1) and silver(II1)ions.

2.2.6 Oxidation of diols Menghani and Bakore7' studied the silver ion-catalysed oxidation of pinacol. Acetone is the main product, and the stoichiometry is presumably

SzOi-

+ (CH,)ZC-C(CH,)z I I OH OH

= 2 (CH,)ZC=O

+ 2 HSO;

(55)

2

O X I D A T I O N S BY P E R O X O D I S U L P H A T E

467

In the absence of silver ions the oxidation is extremely slow. The reaction is firstorder with respect to both peroxodisulphate and silver ions, and the rate is independent of the pinacol concentration. Menghani and Bakore suggest a chain mechanism involving sulphate radical-ions and silver(I1) ions. The same work e r ~ ' ~ ' found ~' the same type of rate equation for the oxidations of propane1,3-dioland butane-1,3-diol, and propose the same mechanism. 2.3

O X I D A T I O N OF A L D E H Y D E S A N D K E T O N E S

Compared with the oxidation of alcohols the oxidation of aldehydes and ketones by peroxodisulphate is generally slow. Formaldehyde and acetaldehyde are oxidised to the corresponding acids, viz. RCHO+SzOi-+H20 = RCOOH+2 HSO,

(56)

Subbaraman and Santappa" studied the oxidations of formaldehyde and acetaldehyde in de-aerated solutions, both in the presence and absence of silver ions. When the concentration of aldehyde is much less than that of peroxodisulphate, the rate equation for reaction in absence of silver ions is

With excess formaldehyde the rate equation becomes

Oxygen inhibits the oxidations of both aldehydes. Subbaraman and Santappa suggest that the hydrated forms of the aldehydes react via chain mechanisms similar to those proposed for alcohols. Edwards et al? in their study of the oxidation of ethanol (section 2.2.2) proposed a mechanism for the oxidation of acetaldehyde involving the non-hydrated form, viz. reactions (42)-(44). Acetone and cyclohexanone are oxidised at an appreciable rate only when silver ions are present. Acetone is oxidised to acetic acid and carbon dioxide, and the process is first-order with respect to both peroxodisulphate and silver ions, and zero-order with respect to acetone (Subbaraman and Santappa", Bekier and Kijowski"). The former workers observed that de-aeration has little effect on the rate, and suggested a chain mechanism involving CH,COCH2 radicals. Cyclohexanone is oxidised about 50 % more rapidly than acetone (Subbaraman and Santappa'l). References p p . 489-492

468

OXIDATION OF ORGANIC COMPOUNDS

2.4

O X I D A T I O N OF C A R B O X Y L I C A C I D S

2.4.1 Oxidation of formic acid The stoichiometry of the oxidation is HCOOH+ SzOi- = COz +2 HSO;

(59)

Srivastava and G h ~ s h 'report ~ that the kinetics are first-order with respect to peroxodisulphate and zero-order with respect to formic acid, but Kappanas4 reports first-order kinetics with respect to each reactant. The effect of trace amounts of metal ions and of oxygen on the rate is uncertain, and discussion of the mechanism is of doubtful significance at present. However, the reported observations definitely indicate a chain mechanism. Thus Srivastava and Ghoshs5 found an induction period in the oxidation, and report that halide ions inhibit the reaction (inhibition by halide ions is a feature of reactions involving hydroxyl radicals). In a study of the silver ion-catalysed oxidation, Gupta and Nigams6 found that the reaction is approximately first-order with respect to both peroxodisulphate and the catalyst, and zero-order with respect to the substrate. The y-ray-initiated oxidation of formic acid by peroxodisulphate has been studied by Hart", who reports the dependence of the yield of carbon dioxide on various factors, but does not give any kinetic data.

2.4.2 Oxidation of oxalic acid The oxidation of oxalic acid has been studied thoroughly by Allen et aI.88-91. The stoichiometry of the reaction is

szo;-+c,o:-

= 2 COZ+2

s0:-

The reaction is very sensitive to metal ion catalysis, particularly by Cuz+ and Ag+, and oxygen inhibits the reaction. Po and Allen" studied the uncatalysed reaction in oxygen-free solutions containing l o w 5M EDTA to ensure that the concentrations of free metal ions were insignificant. Under these conditions the reaction is first order with respect to peroxodisulphate and the rate is essentially independent of oxalate concentration (there is a slight increase in the first-order rate coefficient with increase of oxalate concentration). Ally1 acetate inhibits the reaction and reduces the rate to that observed in the absence of oxalate. In the range pH 0.5-10.3 a rate maximum occurs at pH 4.5. The first-order rate coefficient for the reaction using 0.08 M disodium oxalate is expressed by

k, = 1.61 x 10" exp (-32,7OO/RT) sec-'

2

OXIDATIONS B Y PEROXODISULPHATE

469

The temperature range studied was 45-65 "C. The postulated mechanism is

szo;- 2 2 so, SO, +H,O

OH + C20:-

2 HSO, +OH 2 CO, + CO; + OH-

(62) (63)

It: H,O

(64)

co;+s,o;- r: co2+so,+so:-

(65)

OH-+H+

co;+so, 2co2+so:-

(66)

Provided that inequality (68) is assumed, the stationary-state approximation applied to the concentrations of the radicals leads to the rate equation

in agreement with the first-order kinetics. By equating (k, kz k5k6)* to the observed first-order rate coefficient, and using the appropriate value of k, due to Kolthoff and Millers9 (section 2.1.1), Po and Allen showed that the inequality is justified. An alternative mechanism, proposed by Saxena and Singha19', involving attack on oxalate by sulphate radical-ions rather than hydroxyl radicals is inconsistent with the observed kinetics. From equation (67) it follows that the observed activation energy, E A ,is related to the activation energies of steps (61), (62), (65), and (66) by the equation

El and E , are known (Kolthoff and Millers9, and Kalb and Allengo),and E6 must be very small, so Po and Allen deduced that E, = 24.5 kcal.mole-', indicating that reaction (62) is fairly slow, and that the steady-state concentration of sulphate radical-ions is relatively high. The p H dependence of the reaction rate is only partly explicable on the basis of the proposed mechanism. The rise from pH 0.5 to the maximum at pH 4.5 is probably due to the ionisation of oxalic acid. The inhibition by oxygen is explained by a reversible reaction between oxygen and carbon dioxide radical-ions References p p . 489-492

470

OXIDATION OF O R G A N I C COMPOUNDS

co, + 0 2 z? 02co; and by a new termination step. OZCO,

+ co; + 0 2 +c20:-

(71)

Oxygen is known to inhibit other reactions involving CO; radicals. Early work showed that the rate of the silver ion-catalysed oxidation of oxalate is much faster than the oxidations of other substrates under similar conditions King93). Allen" showed that with solutions of very low copper concentration, but not de-aerated, the rate is only slightly faster compared with other substrates. However, Kalb and Allen" found that oxygen is a powerful inhibitor of the silver ion-catalysed oxidation, and that in the absence of oxygen low concentrations of copper have no effect on the rate. They studied the silver ion-catalysed reaction in the absence of oxygen. With peroxodisulphate concentrations greater than 0.004 M the rate equation is

- d[S20'-1 dt

= k,[S,O~-]*[AgNO,]*

Below pH 0.5 the order with respect to silver changes to first. The rate coefficient

k, increases in the range pH 0-6. From measurements in the temperature range 1%35 "C,and at pH 6, k, is expressed by k, = 5.90 x 10" exp (- 16,4001RT) 1.mole-'.sec-' Ally1 acetate inhibits the reaction, but the maximally inhibited rate is 2.5 times faster than the silver ion-catalysed decomposition of peroxodisulphate in the absence of oxalate. With peroxodisulphate concentrations less than or equal to 0.004 M , the rate becomes proportional to the peroxodisulphate concentration squared and independent of the catalyst concentration, uiz.

During the reaction under the latter conditions a slight precipitate of silver and traces of silver oxalate appear. Kalb and Allen proposed the mechanism Ag'

+ S20'- 5 Ag2++SO; +SO:-

2 SO:-+Ag2+ Ag2++ C20a- 4 Ag' + C 0 2+COT SO; +Ag+

(74) (75) (76)

2

OXIDATIONS B Y PEROXODISULPHATE

co; + szo;- 2 co, + so; + s0:-

2 czo:+Ag+ 2 CO,+Ago

2 co;

CO;

47 1 (65)

(77) (78)

The stationary-state approximation leads to the rate equation (assuming long chains)

where

If the concentration of peroxodisulphate is such that X >> 1, then

Equation (80) agrees with the rate equation observed for [S,Oi-] > 0.004 M . If the peroxodisulphate concentration is low enough to make X << 1, then (1 +X)i approximates to 1 ++X , and the resulting equation

agrees with the observed equation (73). The relative importance of the termination reactions (77) and (78) determines the form of the rate equation. Reaction (77) predominates at the higher peroxodisulphate concentrations, and reaction (78) at the lower. Silver is precipitated in the latter case. From equation (80) it follows that the activation energy E,, observed at higher peroxodisulphate concentrations, is related to the activation energies of reactions (74), (65), and (77) by the equation

E7 is known (Bawn and M a r g e r i ~ o n ~and ~ ) El, is expected to be very small, so Kalb and Allen deduced that E 5 = 7.4 kcal.mole-'. Kalb and Allen" discuss the modifications to the mechanism required to allow for the formation of the complex ion AgC,O; at higher pH values. The variation of rate with pH can be accounted for on this basis. However, the first-order References pp. 489-492

472

OXIDATION OF ORGANIC COMPOUNDS

dependence on the silver ion concentration at low pH could not be accounted for. Sengar and G ~ p t claim a ~ ~that the silver ion-catalysed oxidation is subject to inhibition by oxalate, and they report other features not noticed by Kalb and Allen. Sengar and Gupta do not mention any control of the amounts of copper and oxygen in their solutions, and possibly their results were affected by the presence of these substances. Ben-Zvi and AllenS9 studied the cupric ion-catalysed reaction. As in the uncatalysed reaction the kinetics are first order with respect to peroxodisulphate and zero order with respect to oxalate. The order with respect to total copper is half, the rate expression being

The activation energy is 32.2 kcal.mole-I (from measurements in range 25-44 "C). The rate is not entirely reproducible and is affected by the nature of the vessel surface. The rate is independent of pH in the range pH 3.8-6.3. Oxygen inhibits the reaction slightly, but does not affect the rate law. Ally1 acetate inhibits the reaction completely, as in the uncatalysed reaction. The postulated mechanism is

2 2 so, so, +cu"(C,o,);- 2 cu"'(C,o,); +so:cu"'(c204); 2 c u " c 2 0 , + co, + co, cu"c,04+c,o:2 cu'yc204)2co; + s20;- 1:CO, + so, +so:co, +so; 2co, + so:s,o;-

(61) (84)

(85) (86) (65)

(66)

The stationary-state approximation leads to the correct rate equation (87), below, provided that inequality (88) is assumed. The inequality is justified by the known equilibrium data for the dioxalatocuprate(I1) ion.

Alternative mechanisms which lead to the same rate law are discussed by Ben-

2

OXIDATIONS BY PEROXODISULPHATE

473

Zvi and Allen. From equation (87) it follows that the observed activation energy, E,, ,is expressed by the equation

El and E5 are known (see above) and E6 must be very small, so E l , = 23.5 kcal. mole-' (the value 28.6 kcal given by Kalb and Allen" appears to be incorrect). Hence reaction (84) is quite slow, and the steady-state concentration of sulphate radical-ions relatively high. This would account for reaction (66) being favoured as the chain-termination step, rather than reaction (77), the dominant termination step in the silver ion-catalysed reaction.

2.4.3 Oxidations of acetate and other carboxylate ions yielding products similar to those produced by anodic oxidation The oxidation of acetate by peroxodisulphate is much slower than that of formate. Glasstone and Hickling' showed that the products, which include carbon dioxide, methane, ethane, and ethylene, are similar to those produced by the anodic oxidation of acetate ions (Kolbe electrolysis), and they inferred that the same organic radicals are formed as intermediates. Similar results are reported by Eberson et al.96 for the oxidations of ethyl tert.-butyl-malonate, tert.-butylcyanoacetate, and tert.-butyl-malonamate ions. The oxidations of these ions and of acetate by peroxodisulphate are first order with respect to peroxodisulphate and zero order with respect to the substrate. Mechanisms involving hydroxyl radicals are excluded because the replacement of peroxodisulphate by Fenton's reagent leads to different products, so Eberson et al. infer that the initial attack on the substrate is by sulphate radical-ions. Sengar and P a n d e ~ 'report ~ that the rate of the silver ion-catalysed oxidation of acetate is independent of the peroxodisulphate concentration.

2.4.4 Oxidation of cr-hydroxy acids The oxidations of lactic, malic, and tartaric acids have been studied. In each case carbon dioxide is produced, and in addition lactic acid gives acetaldehyde, and malic acid gives malonic acid. The kinetics are not well defined, and most of the studies show changes of order during a run. In general, the observations are in accord with chain mechanisms. Kumar and Saxena9'* 9 9 found that the oxidation of lactic acid is first-order with respect to peroxodisulphate and zero-order with respect to the acid. Bakore and Joshi"' studied the silver ion-catalysed oxidation, and found approximately first-order kinetics with respect to both References p p . 489- 492

474

O X I D A T I O N O F O R G A N I C COMPOUNDS

peroxodisulphate and silver ion, and zero order with respect to lactic acid. They suggest a chain mechanism involving attack by silver(I1) ions on lactic acid to report produce CH,cHOH radicals. Venkatasubramanian and Sabesan',. that the rate of the reaction in 0.1 M sulphuric acid becomes dependent on the lactic acid concentration below 0.35 M concentration of the latter. The same workers7, report that the oxidation of ethyl lactate is slower than that of the acid, and is first order with respect to peroxodisulphate and zero order with respect to the ester. Kumar and Saxenaio2 studied the oxidation of malic acid (in the absence of silver ions) and found approximately first-order kinetics with respect to peroxodisulphate, and an order varying from zero to first with respect to malic acid. Similar results are reported by Venkatasubramanian and S a b e ~ a n Saxena ~ ~ . and Singhalio3 studied the oxidation of tartaric acid (in absence of silver ions) and found features very similar to those reported for the malic acid oxidation.

2.5

O X I D A T I O N O F N I T R O G E N - C O N T A I N I N G COMPOUNDS, OTHER T H A N AROMATIC AMINES

2.5.1 Oxidation of hydrazobenzene In acetonitrilelwater solvent hydrazobenzene is oxidised to azobenzene (Whalley et al.'04), uiz.

Oxygen was excluded because it reacts directly with hydrazobenzene in acetonitrile/ water solvent. The kinetics are approximately second-order (first-order with respect to each reactant). The second-order rate coefficients increase with decreasing initial concentration of substrate, and with increasing initial concentration of peroxodisulphate. With equal concentrations of the reactants below 0.006 M the second-order rate coefficient is almost independent of the initial concentrations. Whalley et al. suggested the mechanism

2

OXIDATIONS B Y PEROXODISULPHATE

415

The rate equation derived from this mechanism is in accord with most of the observed features, but it predicts that with excess substrate the second-order rate coefficient should decrease during a run, whereas the observed rate coefficient always increases during a run, irrespective of whichever reactant is in excess. Whalley et al. suggest that incomplete dissociation of peroxodisulphate in the solvent might be responsible for the discrepancy. Another discrepancy is pointed out by Wilmarth and Haim57, but these authors agree with Whalley et aZ. in concluding that the initiation step is reaction (91) rather than the spontaneous fission of the peroxodisulphate ion.

2.5.2 Oxidation of formamide and acetamide Agrawal and Mushran1'' studied the kinetics of the silver ion-catalysed oxidation of acetamide. The stoichiometry is uncertain, but acetic acid and nitrogen are the main products. The rate is approximately first-order with respect to both peroxodisulphate and silver ions, and is almost independent of the substrate concentration. No definite conclusions regarding the mechanism can be drawn, but the kinetics suggest a chain process. Agrawal et aZ.lo6 report similar results for the oxidation of formamide.

2.5.3 Oxidation of aminoalcohols Beileryan et aZ.'07 have studied the kinetics of oxidation of several N, Ndiethyl-aminoalcohols. In the presence of oxygen and with excess substrate, the reactions are three-halves order with respect to peroxodisulphate, and first-order with respect to the substrate. A chain mechanism is proposed.

2.6

O X I D A T I O N OF S U L P H U R - C O N T A I N I N G C O M P O U N D S

Mercaptans are oxidised to disulphides by peroxodisulphate. Eager and Winkler"' studied the kinetics of the oxidations of n-butyl, n-octyl, and ndodecyl mercaptans in acetic acid/water solvent (80 ml acid+20 ml water) and found first-order kinetics with respect to peroxodisulphate. The first-order rate coefficient increases with increase of mercaptan concentration, and reaches a limit at about 5 x M mercaptan. A decrease in the first-order rate coefficient with increase of the initial peroxodisulphate concentration was observed and attributed to a salt effect. Eager and Winkler suggested a mechanism involving sulphate radical ions. Levitt' O 9 proposed a mechanism involving sulphur tetroxide, but there is no evidence for its formation in peroxodisulphate oxidations. References pp. 489492

476

OXIDATION OF ORGANIC COMPOUNDS

Diphenyl sulphoxide, thiodiglycol sulphoxide, and diethyl sulphoxide are oxidised to the corresponding sulphones. Howard and Levitt" studied the kinetics of oxidation of these compounds at pH 8 in a phosphate buffer (aqueous solution) and found the rates to be first-order with respect to peroxodisulphate and independent of the substrate concentration over the limited range employed (0.01-0.02 M ) . The removal of oxygen had no effect on the rate. Diethyl sulphide is oxidised very rapidly initially, then at the same rate as diethyl sulphoxide. Howard and Levitt concluded that the sulphide is first oxidised to the sulphoxide which in turn is oxidised to the sulphone, but Wilmarth and Haim" point out that this interpretation cannot be correct, and conclude that the reaction must be more complex. Douglas"' studied the oxidation of thiophenolate ion

'

C,H5S- + 2 SzOi- = (C6H5S)z+2SO:-

(95)

The reaction is approximately first-order with respect to each reactant (the secondorder rate coefficient increases with increase of substrate concentration), and catalysis by hydroxide ions is observed. Henderson and Winkler' l 2 studied the ferrous ion-catalysed oxidation of thioglycolicacid to dithioglycolic acid. The rate is sensitive to traces of metal ions, and reproducible results could not be obtained in the absence of the catalyst. The oxidation is first-order with respect to both peroxodisulphate and ferrous ions, and zero-order with respect to the substrate. The second-order rate coefficient is approximately equal to that determined in the absence of the substrate, so Henderson and Winkler suggested that the ratedetermining step is the oxidation of ferrous to ferric ions, as in reaction (96), and that this is followed by reaction (97) and then rapid oxidation of thioglycolic acid by ferric ions.

2.7

OXIDATIONS OF PHENOLS A N D AROMATIC AMINES

As noted in the introduction, the oxidations of phenols and aromatic amines by peroxodisulphate proceed via ionic rather than free-radical mechanisms.

2.7.1 Oxidation of phenols (Elbs reaction) In alkaline solution phenols are readily oxidised by peroxodisulphate. The normal result of the reaction is the substitution of sulphate in the position para to the hydroxy group, viz.

OXIDATIONS BY PEROXODISULHPHA’I E

I

477

_

oso3

Some ortho substitution usually occurs as well. For preparative purposes the sulphate is hydrolysed by acid to give a para-diphenol. Behrman and Walker113 studied the kinetics of oxidation of 2-hydroxy-pyridine and o-nitrophenol. For both these compounds the kinetics are second order, uiz.

- dCS2” dt

= k,[S,Oi-][phenol]

(99)

The rate coefficient k, is proportional to the concentration of phenolate ion; thus for 2-hydroxy-pyridine (pK, = 11.6) oxidation below pH 8 is slower than the spontaneous decomposition of peroxodisulphate. For the reaction of 2-hydroxypyridine in 2 M sodium hydroxide, the variation of k, with temperature is expressed by

k,

=

7.56 x lo8 exp (- 16,30O/RT) 1.mole-’.sec-’

(temperature range studied = 18-40 “C).The reaction is not catalysed by cupric or ferrous ions, nor is it inhibited by oxygen, allyl acetate, or allyl alcohol (Behrman1I4). Thus all of the observations point to an ionic rather than a free-radical mechanism. Electron-releasing substituents increase the rate, and the reaction is considered to be an example of nucleophilic displacement on peroxide oxygen. Two general reaction schemes, (a) and (b), are possible in view of the ambident nature of the nucleophile, viz. 0-

OOSOJ

I

SCHEME (a)

References pp. 489-492

I

478

O X I D A T I O N OF O R G A N I C COMPOUNDS

+

so ,-;

0

- $g H

SCHEME (b)

+

'-8 so',-

OSO;

+ H C

oso;

Scheme (a) involves rate-determining attack by the peroxodisulphate at oxygen, and (b) involves rate-determining attack at carbon. In an attempt to distinguish between (a) and (b), Behrman'l4 obtained kinetic data for a series of 39 phenols, and plotted log k , against Hammett'o constants in two ways corresponding to the two schemes. Thus, (i) For attack at oxygen, log k , should correlate with o: for m-substituted phenols, and with op' for o-substituted phenols. (ii) For attack at carbon, log k , should correlate with 0, for 0- and p-substituted phenols, but the correlation of m-substituted phenols with op' should be poor due to ortho effects. The best overall correlation is obtained by plotting according to (i), but Behrman concluded that Hammett plots are not of great value in the present case because peroxodisulphate is not a typical electrophilic reagent, also AS' is not constant for the series of phenols (however, its variation is related in an approximately linear manner to the variation of AH'). Probably there is a varying contribution of attack at oxygen and attack at carbon. Kinetic data obtained for some 2,6 and 2,4 disubstituted phenols show that crowding round the phenolic oxygen accelerates the rate when electron-releasing substituents are used. In these cases attack must be at carbon. The values of A H f for the series of phenols studied by Behrman range from 11.7 to 15.8 kcal.mole-I, and of AS' from -12 to -30 cal.deg-'.mole-l. These activation parameters are similar to those found for other nucleophilic displacements on peroxide oxygen (Edwards'' ', and Behrman' l').

2.7.2 Oxidation of aromatic amines (Boyland-Sims reaction)

This oxidation is similar to the Elbs reaction, except that ortho substitution of sulphate normally occurs, e.g.

2

479

OXIDATIONS B Y PEROXODISULPHATE

The yield of o-amino aryl sulphate is generally 10-40 %. If both ortho positions are blocked then para substitution results. For preparative purposes the sulphate is hydrolysed by acid, as with the Elbs reaction. Behrman1I7 studied the oxidation of 2-amino-pyridine,and showed that the kinetics are second -order for the major part of reaction, viz.

-’

- d[S2 ” dt

=

k,[S2 Oi-1 [amine]

The pH dependence of the rate shows that the free amine is the reactive species. For the reaction in 0.475 A4 potassium hydroxide, k2 is expressed by

k2

=

4.55 x lo7 exp (15,70O/RT)1.mole-’.sec-’

The salt effect is small and positive, consistent with the attack of an ion on a neutral molecule. Ally1 acetate has no effect on the rate or yield of product. Electron-releasing substituents accelerate the reaction. Thus the features of the reaction are similar to those of the Elbs reaction, and point to a mechanism involving nucleophilic attack of the amine on peroxide oxygen. In order to determine whether attack occurred by nitrogen or carbon, Behrman’ compared the rates of oxidation of amines such as 2-amino-5-methyl-pyridine and 2-amino6-methyl-pyridine,and found that in all cases the order of reactivity is consistent with attack by nitrogen. Furthermore, the rate of oxidation of 3,5-dideutero2-amino-pyridine is the same as for the normal material, and the yield of product is the same. Behrman proposed the reaction scheme

’

+

amine S 2 0 i-

ki

+

X

ki

+

o-amino aryl sulphate

ka

X + n SzOi- + further oxidation products+NH, The intermediate X, present in low concentration, is probably an aryl hydroxylarnine-U-sulphonate.The further oxidation products consist mainly of polymeric humic acid. The stationary-state approximation applied to [XI shows that the ratio of the rate of formation of o-amino aryl sulphate to the rate of formation of other products is kJk2[S20g - I”, in qualitative agreement with the dependence of the yield of o-amino aryl sulphate on peroxodisulphate concentration, and the deviation from first-order kinetics with respect to peroxodisulphate towards the end of a run.

References pp. 489-492

480

O X I D A T I O N OF O R G A N I C C O M P O U N D S

3. Kinetics of oxidations by peroxomonosulphate 3.1

INTRODUCTION

Peroxomonosulphuric acid (Caro's acid), H,S05, has the structure (1). 0 II

HO-S-OOH II

0 (1) One of the protons is highly acidic, but the other (attached to the peroxide group) is weakly acidic, and the second pK, is 9.4 (Ball and Edwards"'). At room temperature in aqueous solutions of about pH 9 the spontaneous decomposition of the acid

HSO; = HSO;+$O, is rapid. A small amount of peroxodisulphate is also produced (Goodman and Robson" '). The decomposition is second-order with respect to the total peroxomonosulphate species, the rate reaching a maximum at pH 9.5 (Ball and Edwards"'). In contrast to the spontaneous decomposition of peroxodisulphate no radicals are involved, and the rate-determining step is considered to be the attack of the dianion, SO:-, on the monoanion, HSO; (Ball and Edwards"', Goodman and Robson' '). Certain metal ions catalyse the decomposition very effectively (Ball and Edwards'").

3.2

OXIDATIONS

Peroxomonosulphate oxidises a wide variety of organic compounds, but very few kinetic studies have been carried out. All of the available evidence points to ionic rather than free-radical mechanisms.

3.2.1 Oxidation of aromatic amines and nitroso compounds Ogata and Tabushi" ' studied the kinetics of oxidation of some N,N-dimethylanilines in aqueous solutions of pH 1-12 at 25 "C (some runs were performed at other temperatures in the range 15-30 "C). The product of oxidation is an amine oxide, viz.

3

OXIDATIONS BY PEROXOMONOSULPHATE

Ph(Me),N+ HSO,

f

=

Ph(Me),N-0-

48 1

+ HSOT

The kinetics are second-order (first with respect to each reactant). The variation of rate with pH shows a plateau in the range ca. pH 5-9, with a fall in rate on each side of the plateau. The rate is unaffected by the addition of small amounts of hydroquinone or benzoyl peroxide. Ogata and Tabushi accounted for the pH dependence by assuming the rate-determining step to be the attack of undissociated peroxomonosulphuric acid on the free amine. However, they based their discussion on a misconception, viz. that the$rst pK, of the acid is about 9.5, and that the dissociation of the second proton is inappreciable. The pH dependence is correctly accounted for the assuming that attack of the monoanion, HSO;, on the free amine is rate-determining (there is a kinetically equivalent alternative, viz. attack of the dianion, SO:-, on the protonated amine, but this is a very unlikely alternative because there is evidence for the amine acting as a nucleophile, also in acid solution the concentration of the dianion is very low). The rate coefficients at 25 "C for attack of HSO; on the free amines are 50, 63, 108, and 124 1.mole- .sec-' for the oxidations of p-chloro-, unsubstituted, p-methyl-, and pmethoxy- N,N-dimethylanilines respectively. Thus the order of reactivity is consistent with the amines acting as nucleophiles. With acid solutions the order of reactivity is reversed, because electron-releasing substituents increase the basicity of an amine, and therefore reduce the concentration of free amine in acid solution. The reaction probably involves nucleophilic displacement on the outermost peroxide oxygen. Ibne-Rasa and EdwardsI2' have put forward arguments for believing that a solvent molecule is involved in the transition state of this and similar oxidations by organic peroxy acids, viz.

(The bonds in HSO; which are breaking in the transition state are represented by - - -). Gragerov and LevitlZ3 studied the oxidations of some aromatic amines and nitroso compounds by peroxomonosulphate in H2180, and found that the products were unlabelled. Therefore the formation of hydroxyl radicals (by attack of peroxomonosulphate on the solvent) and their participation in the oxidations is excluded. In a comparison of the rates of oxidation of nitrosobenzene by various peroxy References pp. 489-492

482

O X I D A T I O N OF O R G A N I C C O M P O U N D S

acids, Ibne-Rasa et al.124quote a second-order rate coefficient of 3.5 x I.mole-'.sec-' for the oxidation by peroxomonosulphate at 30 "C in 47 % ethanol-water solvent. From a comparison of the rate coefficient with that observed for the oxidation of nitrosobenzene by peroxoacetic and peroxochloracetic acids, they concluded that the mechanism is similar to that described above for the oxidation of N,N-dimethylanilines.

3.2.2 Oxidation of diphenyl sulphide Gragerov and Levit',' studied the oxidation of diphenyl sulphide by peroxomonosulphate in ethanol-acetic acid solvent containing H2180, and found that the products, diphenyl sulphoxide and diphenyl sulphone, were unlabelled. They concluded that the oxidation is heterolytic and does not involve free-radicals.

4. Kinetics of oxidations by hypochlorous acid

4.1

INTRODUCTION

In aqueous solution hypochlorous acid is in equilibrium with molecular chlorine, viz. C1,

+ H 2 0 + H + + C1- + HClO

(1)

A low equilibrium concentration of chlorine monoxide is also present 2 HClO

+ C1,O + H,O

(2)

at 27 "C (Lister'26), and The equilibrium constant for reaction (1) is 5.5 x for reaction (2) 9.6 x at 0 "C (Go ld ~c h mid t'~ ~Hypochlorous ). acid is a weak acid having a pK, of 7.49 at 20 "C (Shilov et Oxidations by hypochlorous acid are usually studied in alkaline solution to minimise competing reactions of molecular chlorine. The oxidations of many organic substrates, e.g. amines and phenols, by hypochlorous acid are very complex and often give rise to a large number of products. The few kinetic studies that have been carried out are mostly confined to reactions whose stoichiometry is fairly well defined, and the general conclusion is that the rate-determining step is the attack of HClO on the substrate. However, for the oxidation of formate ions (next section) one group of workers conclude that chlorine monoxide is the active species. There is no conclusive evidence for the anion C10- being the active species in any oxidation.

4

OXIDATIONS BY HYPOCHLOROUS ACID

4.2

483

OXIDATION OF FORMATE

Shilov et ul.12’ studied the rate of oxidation of formate ions in phosphate and carbonate buffers, and showed that the reaction with molecular chlorine is negligible in solutions of pH > 6. At 20 “C the rate of reaction with hypochlorous acid is constant in the range pH 5.5-7, then it decreases with increase of pH, and becomes negligible at p H 13. The kinetics are second-order with respect to hypochlorous acid, and first with respect to formate ions. In alkaline solution hydroxide ion catalysis is apparent; viz.

- d[Cl= k2 [HClo]2[HCOO-] dt

+ k2’[HC10]2[HCOO-l[OH-l

(3)

where C = “active chlorine”. The second-order kinetics with respect to hypochlorous acid were taken to indicate that chlorine monoxide is the active species. In contrast to rate equation (3), Lister and Rosenblum’ 2 9 report that the reaction in unbuffered alkaline solutions is first-order with respect to both hypochlorite and formate ions, uiz.

= k’[HCIO][HCOO-]

(5)

Lister and Rosenblum suggest that the rate-determining step is the attack of HClO on HCOO-, uiz.

HClO + HCOO-

-+

CI- + H2C0,

11

2 H++CO:-

From rate measurements at 60 and 70 “C, the activation energy for reaction (6) is 6.7 kcal.mole-’, and the A-factor is 3.4 x lo3 1.mole-’.sec-’. The latter is unusually small, but Lister and Rosenblum point out that similar low values have been obtained for the oxidations of bromide and nitrite ions by hypochlorous acid. They suggest that the formation of the transition state requires a special orientation of the reactant molecules.

References pp. 489-492

484

OXIDATION O F O R G A N I C COMPOUNDS

4.3

M I S C E L L A N EOU S OX I D AT1 O N S

4.3.1 Cyanate

The oxidation of cyanate by hypochlorous acid is a complex reaction. According to Lister130,the stoichiometry of the main reaction is given by equation (7), and that the most plausible secondary reaction is represented by (8), uiz. 3 C 1 0 - + 2 C N O - + H 2 0 = 2HCO;+3C1-+N2

(7)

4 C10- +CNO- +2 OH- = COi- +4 CI- +NO; + H 2 0

(8 )

Lister studied the kinetics in 0.2-1.0 M sodium hydroxide, and deduced the following rate equation for the main reaction

- d [C10 -1 - k [ (210-1[CN 0-1 dt

L-OH-]

(9)

Lister suggested that the rate-determining step is a bimolecular reaction between hypochlorous acid and cyanate ions. The subsequent steps are very uncertain. From measurements in the temperature range 50-75 "C, the activation energy of the rate determining step is 15.1 kcal.mole-', and the A-factor is 3.2 x 10" 1. moIe-'.sec-'.

4.3.2 Glycolaldehyde and glucose For both these substrates Shilov and Y a ~ n i k o v ' ~showed ' that hypochlorous acid is the active species in neutral or weakly alkaline solutions. In the range pH 5-9.5 the rate equation is

-_ dCcl_--k[C][S][OH-] dt where C = "active chlorine", and S = substrate. Above pH 10 the reaction becomes zero-order with respect to hypochlorite, indicating that the formation of the enolate ion of the substrate is rate determining. Grigor et ~ 1 . lstudied ~ ~ the oxidation of glucose in a phosphate buffer at pH 7.0, and observed three distinct parts of the glucose concentration and optical rotation versus time curves. They deduced that the order of reactivity of the forms of glucose is aldehyde form > a-glucose > P-glucose.

5

OXIDATIONS B Y H Y P O B R O M O U S A C I D

485

4.3.3 Hydroxy acids The oxidations of lactic acid, d-gluconic, and pyruvic acids have been studied by Shilov and Y a ~ n i k o v ' ~They ~ . found that the rate of reaction in each case is very low in acid solution, and that the rate increases with increase of pH to a maximum a t pH 7. In solutions of pH 7 the kinetics are first-order with respect to both "active chlorine" and the substrate. Shilov and Yasnikov deduced that hypochlorous acid is the principal active species, and that the mechanism involves the formation of an ester of hypochlorous acid as the first step, followed by elimination of hydrochloric acid and the formation of a carbonyl group.

-

4.3.4 Benzaldehyde sulphonates Kozinenko and S h i 1 0 v ' ~studied ~ the kinetics of oxidation of the 0-,m-,and p benzaldehyde sulphonate ions. In buffered solutions of pH 4-13 the rate equation is

where C = "active chlorine", S = substrate, and B = buffer anion. The buffer catalysis term is explained by nucleophilic addition of B to the -CHO group, and the subsequent attack of hypochlorous acid at the hydrogen atom. 5. Kinetics of oxidations by hypobromous acid

5.1

INTRODUCTION

In aqueous solution hypobromous acid is in equilibrium with molecular bromine, viz. Br, + H 2 0 f H + +Br-

+HBrO

(1)

The equilibrium constant for reaction (1) is 9.6 x l o p 9 at 25 "C (Pink'35). Thus the proportion of bromine present as hypobromous acid is very small in solutions of pH < 6. Hypobromous acid is a weak acid, having a pKA of 8.66 at 25 "C (Flis et On account of the equilibrium (l), reactions of hypobromous acid are frequently accompanied by reactions of molecular bromine, and the contribution of any reaction due to the acid has to be determined from the effects of changing the pH and the addition or removal of bromide ions. Perlmutter-Hayman et al.137-9have shown that with the majority of organic substrates, oxidaReferences p p , 489-492

486

O X I D A T I O N OF O R G A N I C C O M P O U N D S

tion by molecular bromine is much faster than oxidation by hypobromous acid. The only known exception is the oxidation of oxalic acid (next section). For the substrates ethanol, acetaldehyde, and glucose the contribution of hypobromous acid oxidation has been determined (section 5.3), and for 2-propanol an upper limit has been placed on the rate coefficient for oxidation by hypobromous acid (Perlmutter-Hayman and W e i ~ s m a n l ~In ~ ) all . examples studied the kinetics of oxidation are second-order (first-order with respect to both HBrO and the substrate), and it is generally assumed that the rate-determining step is the attack of HBrO on the substrate. There is no conclusive evidence for the anion BrO- being the active species in any oxidation. 5.2

OXlDATION OF OXALIC ACID

Early work on the kinetics of oxidation of oxalic acid by hypobromous acid is reviewed by Griffith et aLi40. The latter workers showed that in solutions of pH < 6 the oxidising agent is HBrO and that no appreciable oxidation by molecular bromine occurs. The rate equation is

- d[HBrol dt

= k[HBrO][HC,O,]

Accordingly, Griffith et al. assumed that the rate-determining step is the bimolecular reaction of HBrO and HC204-, viz. HBrO + H C z O i

+2

CO,

+ H,O + Br-

(3)

Further work carried out by Makower and Liebhaf~ky'~'showed that the rate coefficient k , calculated assuming equation (2), is constant below pH 6.2 but increases in more alkaline solutions. They attributed the increase to a second reaction which could be either BrO-

+ HCzO;

+ COz + Br-

+

HCO;

-+

HCO;+CO,+Br-

(4)

or HBrO+C,OZ-

(5)

From the temperature dependence of the rate below pH 6.2, Makower and Liebhafsky express the rate coefficient for reaction (3) by

k = 1.0 x 10l6 exp (- 15,700IRT)l.mole-'.sec-'

5

OXIDATIONS BY HYPOBROMOUS ACID

487

H i n ~ h e l w o o d 'suggested ~~ that the bromine cation Br' is the active species rather than HBrO. Attack by Br' on the dianion C,O?- is kinetically indistinguishable from attack by HBrO on HC,O,, but Knoller and Perlrn~tter-Hayman'~~ point out that because the concentration of Br' i n aqueous solution is very low, the rate of attack of Br' on C,O:- would have to be considerably greater than the limiting rate imposed by diffusion. They conclude that HBrO is the active species, in agreement with the earlier workers, but they consider that undissociated oxalic acid is attacked as well as the anion HC,O,. They give values of 2.22 x 10' and 2.82 x 10' I.mole-'.sec-' for the rate coefficients for the attack of HBrO on H2C204 and H C 2 0 4 respectively at 20°C. They state that hypobromous acid does not attack the dianion, but their conditions were restricted to the range pH 0.8-5, and they do not comment on Makower and Liebhafsky's evidence for reactions (4) or (5).

5.3

MISCELLANEOUS OXIDATIONS

5.3.1 Oxidation of ethanol and acetaldehyde

Ethanol is oxidised to acetaldehyde, which in turn is oxidised to acetic acid. Perlmutter-Hayman and W e i ~ s m a n n ' showed ~~ that oxidation by molecular bromine predominates. Thus, for ethanol at pH 1.0 and 0 "C they found kSr2= 4x I.mole-'.sec-', whereas kHBrO is about 30 times smaller, and for acetaldehyde under the same conditions, kBr2= 2.1 x l.mole-'.sec-' whereas kHBrO is about 50 times smaller. The reactions with hypobromous acidaremuch more light-sensitive than those with molecular bromine.

5.3.2 Oxidation of glucose The initial product is gluconic acid lactone. Perlmutter-Hayman and P e r ~ k y ' ~ ' showed that at pH 1.25 and 0 "C, the rate coefficient for attack of hypobromous acid on glucose is 2.04 x lo-, I.mole-'.sec-' (the corresponding rate coefficient for attack by molecular bromine is 7 times greater). The rate coefficients both for attack by HBrO and Br, increase with increase of pH beyond 3. This increase is attributed to the much greater reactivity of the glucose anion.

References p p . 489-492

488

OXIDATION OF ORGANIC COMPOUNDS

6. Kinetics of oxidations by hypoiodous acid 6.1 I N T R O D U C T I O N

The equilibrium constant for the hydrolysis of iodine

T,+H,O+

(1 1

H++I-+HIO

is 3 x at 25 "C (Bray and C ~ n o l l y ' ~ consequently ~), the concentration of hypoiodous acid in aqueous solution is very low unless iodide ions are removed or the pH is high. Hypoiodous acid is much less stable than hypobromous and hypochlorous acids, and the study of its properties is consequently difficult. Literature values of its pK, vary from 9.7 to 12.3; according to a more recent determination (Chia'45) it is 10.64 at 25 "C.The few kinetic studies of oxidations by hypoiodous acid lead to similar conclusions to those found for oxidations by the other hypohalous acids, viz. the acid itself is the active species in a bimolecular reaction with the substrate.

6.2

OXIDATIONS

6.2.1 Oxalic and formic acids In the reaction of oxalic acid with iodate, the oxidising species are molecular iodine and hypoiodous acid (principally the latter). The induction period observed in the reaction is due to build-up of the oxidising species; it is shortened by exposure to light or by the presence of metal ions. As soon as the iodine concentration is appreciable, the reaction rate accelerates rapidly on account of the sequence of reactions 12+Hz0

+H++I-+HIO

(COOH), + H I 0 = 2 C 0 2+ H + + I -

(1)

+ HzO

(2)

Abel and Hilferding14h*'47studied the kinetics of the reaction, and showed that the oxidations of oxalic acid and its monoanion by hypoiodous acid are secondorder. The reaction of formic acid with iodate is similar to the reaction of oxalic acid. Abel and Bildermann' 48 studied the kinetics, and showed that molecular iodine and hypoiodous acid are the oxidising species.

7

OXIDATIONS B Y NITROUS ACID A N D SELENIOUS ACID

489

6.2.2 Aldoses Ingles and Israel149 studied the oxidation of several aldoses by hypoiodous acid. The aldehyde group is oxidised to the acid, viz. RCHO+HIO

=

RCOOH+H++I-

(4)

Glucose, for example, gives gluconic acid. The OH groups are not attacked. Ingles and Israel showed that the rate of oxidation reaches a maximum at pH 11.3, and exactly parallels the variation in the hypoiodous acid concentration as the pH is changed. The kinetics were found to be second-order, after the concurrent decomposition of hypoiodous acid had been allowed for.

7. Oxidations by nitrous acid and selenious acid Although the oxidations do not appear to involve the anions of the acids it is convenient to mention them here. Bunton et aZ.1509151 studied the oxidation of ascorbic acid and similar compounds and concluded that the active species were N203 and H,NO:, while Longstoff and Singer' 5 2 suggested that the oxidation of formic acid in nitric acid (no reaction in the absence of nitrous acid) involved NO+ and HNO,. The oxidation of ketones by selenious acid has been investigated by Corey and Schaefer' 53 and by Schaefer' 54. They report that reaction is due to H,Se03 and H3Se03+and that HSe0,- is definitely inactive. REFERENCES 1 E. L. JACKSON, Org. Reactions, 2 (1944) 341. Aduan. Carbohydrafe Chem., 11 (1956) 1. 2 J. M. BOBBITT, 3 G. DRYHURST, Periodate Oxidation of Diol and Other Functional Groups, Pergamon, London, 1970. 4 B. SKLARZ, Quart. Reo. (London), 21 (1967) 3. 5 C. A. BUNTONin Oxidation in Organic Chemistry, Part A, K. B. WIBERG(Ed.), Academic Press, New York, 1965, Chap. VI. 6 L. MALAPRADE, Bull. SOC.Chim. France, 43 (1928) 683. 7 A. K. QURESHI AND B. SKLARZ, J . Chem. SOC.,C (1966) 412. 8 G . J. BUIST,W. C. P. HIPPERSON AND J. D. LEWIS, J. Chem. Soc., A (1969) 307. 9 C. E. CROUTHAMEL, H. V. MEEK,D. S. MARTINAND C. V. BANKS,J. Am. Chem. SOC., 71 (1948) 3031. A. M. HAYESAND D. S. MARTIN,J. Am. Chem. SOC.,73 (1951) 82. 10 C. E. CROUTHAMEL, 1 1 R. CRIEGEE, Sitzber. Ges. Befoerder. Ges. Naturw. Marburg, 69 (1934) 25; Chem. Abstr., 29 (1935) 6820. 12 G. J. BUIST,C. A. BUNTONAND J. H. MILES,J . Chem. Soc., (1957) 4575. 13 G. J. BUISTAND C. A. BUNTON,unpublished work. 14 L. MALAPRADE, Bull. Soc. Chim. France, 1 (1934) 833.

490

O X I D A T I O N O F O R G A N I C COMPOUNDS

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