effect of metal ions on the autoxidation of benzaldehyde.

 Title Page
 Table of Contents
 List of Tables
 List of Figures
 Experimental procedure
 Biographical sketch
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Material Information

Title: effect of metal ions on the autoxidation of benzaldehyde.
Series Title: effect of metal ions on the autoxidation of benzaldehyde.
Physical Description: Book
Creator: Kresge, Edward Nathan,

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Material Information

Title: effect of metal ions on the autoxidation of benzaldehyde.
Series Title: effect of metal ions on the autoxidation of benzaldehyde.
Physical Description: Book
Creator: Kresge, Edward Nathan,

Record Information

Source Institution: University of Florida
Holding Location: University of Florida
Rights Management: All rights reserved by the source institution and holding location.
Resource Identifier: alephbibnum - 000421880
oclc - 11020962
System ID: UF00091322:00001

Table of Contents
    Title Page
        Page i
        Page ii
    Table of Contents
        Page iii
        Page iv
    List of Tables
        Page v
    List of Figures
        Page vi
        Page vii
        Page 1
        Page 2
        Page 3
        Page 4
        Page 5
        Page 6
        Page 7
        Page 8
        Page 9
        Page 10
        Page 11
        Page 12
        Page 13
        Page 14
        Page 15
        Page 16
        Page 17
    Experimental procedure
        Page 18
        Page 19
        Page 20
        Page 21
        Page 22
        Page 23
        Page 24
        Page 25
        Page 26
        Page 27
        Page 28
        Page 29
        Page 30
        Page 31
        Page 32
        Page 33
        Page 34
        Page 35
        Page 36
        Page 37
        Page 38
        Page 39
        Page 40
        Page 41
        Page 42
        Page 43
        Page 44
        Page 45
        Page 46
        Page 47
        Page 48
        Page 49
        Page 50
        Page 51
        Page 52
        Page 53
        Page 54
        Page 55
        Page 56
        Page 57
        Page 58
        Page 59
        Page 60
        Page 61
        Page 62
        Page 63
        Page 64
        Page 65
        Page 66
        Page 67
        Page 68
        Page 69
        Page 70
        Page 71
        Page 72
        Page 73
        Page 74
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        Page 80
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        Page 86
        Page 87
        Page 88
        Page 89
        Page 90
        Page 91
        Page 92
        Page 93
        Page 94
        Page 95
        Page 96
        Page 97
        Page 98
        Page 99
        Page 100
        Page 101
        Page 102
        Page 103
        Page 104
        Page 105
        Page 106
        Page 107
        Page 108
        Page 109
        Page 110
        Page 111
        Page 112
        Page 113
    Biographical sketch
        Page 114
        Page 115
Full Text






August, 1961


The author wishes to express his sincere appreciation to his advisor

and teacher, Dr. W. T. Lippincott, for his empathy and friendship and for

his inspiration and direction of the research.

The author also wishes to express his appreciation to the members of

the Department of Chemistry and to his friends for their encouragement and







I. Aims and Scope 1

II Autoxidation of Benzaldehyde 4


I, General Statement of Experimental Procedure 18

IH Purification of Benzaldehyde 19

I. Preparation and Purification of Other Chemicals 21

IV. Oxidation Apparatus 25

V. Typical Autoxidation Run 29

VI. Analytical Methods 33


I, Manganese Systems 35

I. The Catalytic Decomposition of Perbenzoic Acid 70

II Cobalt Systems 75

IV, Cerium Systems 81



I, Manganese System 88

HL Cobalt System 99

m c Cerium System 102






1. Activation Energies for the Autoxidation of Benzaldehyde 7

2. Effect of Manganese (II) Concentration on the Oxidation o1
Benzaldehyde 37

3. Effect of Manganese (III) Concentration on the Oxidation of
Bonzaldehyde 43

4. Effect of Perbenzoic Acid on the Oxidation Rate at Low
Catalyst Concentrations 45

5. Effect of Initial Benzaldehyde Concentration on the Oxida-
tion Rate 49

6. Effect of Oxygen Pressure on the Oxidation Rate 52

7. Effect of Water on the Manganese Catalyzed Autoxidation
of Benzaldehyde 53

8. Effect of Benzoic Acid on the Rate of Oxidation 56

9. Change in the oxidation Rate with Temperature 61

10. Effect of Manganese (II) with Acetic Anhydride 64

11. Variations in Bonzaldehyde and Perbenzoic Acid Concen-
trations During the Oxidation 65

12. Effect of Cobalt Concentration on the Rate of Benzaldehyde
Oxidation 76

13. Effect of Benzaldehyde Concentration on Oxidation Rate 79

14. Effect of Cerium on the Oxidation of Benzaldehyde 83


1. Oxidation Apparatus 26

2. Calibration of Recorder 28

3. Typical Autoxidation Run 31

4. Determination of the Induction Period 32

5. The Effect of Manganese (II) Ion on the Oxidation Rate of
Benzaldehyde 38

6. Induction Periods for Manganese Catalyzed Autoxidation
of Benzaldehyde 39

7. The Effect of Manganese (HI) Ion on the Extent of Reaction 40

8. The Effect of Manganese (II) Ion at Low Concentrations on
the Rate of Oxidation 42

9. Plot Showing the One-Half Order Dependency on Perben-
zoic Acid 47

10. The Effect of the Initial Benzaldehyde Concentration on
the Oxidation Rate 50

11. Plot of log -dO2/dt versus log [BzH] 51

12. The Effect of Water on the Extent of Reaction 54

13. Perbenzoic Acid Concentrations During the Autoxidation
at Various Catalyst Concentrations 58

14. Increase in the Peracid Concentration During the Induction
Period 59

15. Exponential Build-Up of Peracid During the Induction Period 60


16. Arrhenius Plot for Manganese-Catalyzed Autoxidation 62

17. Autoxidation of Benzaldehyde 66

18. Decomposition of Perbenzoic Acid by Manganese (II)
Acetate 72

19. Decomposition of Perbenzoic Acid by Manganese (III)
Acetate 73

20. Comparison of Manganese (I) and Manganese (III)
Catalyzed Decomposition of Perbenzoic Acid 74

21. The Effect of the Cobalt Concentration on the Benzalde-
hyde Oxidation Rate 77

22. Variation of the Oxidation Rate with Benzaldehyde Con-
centration 80

23. Typical Cerium Catalyzed Run 82

24. The Effect of Cerium Ions on the Induction Period 85

25. The Effect of Cerium Ions on the Post Induction Oxida-
tion Rate 86


I. Aims and Scope

Transition metal ions have been used as catalysts for many types of

free radical reactions. The nature of their role in these reactions has not

been fully established, although considerable effort has been expended to

elucidate the detailed mechanistic involvement of the metal ions in the reac-

tion sequence. The autoxidation of alcohols, phenols, ethers, amines, enols

and hetols is catalyzed by heavy metal ions, but the most extensively studied

reactions are the autoxidations of hydrocarbons and aldehydes.*

The autoxidation of aldehydes proceeds with ease, even at room

temperature, and undoubtedly proceeds by the chain reaction sequence:

RCHO + In' -- InH + RCO- (1)

RCO. + 02 ----RCOOO (2)

RCOOO + RCHO --*-RCO + CO + OOH (3)

where In* is an initiator radical. The peracid that is produced by this reac-

tion then reacts with unreacted aldehyde to yield the final product, an acid, as

indicated by the equation


*References and a summary of this work are given in Part H of the


To be effective as a catalyst the metal ion must have two readily ac-

cessible oxidation states differing by one unit. Examples are cobalt (II) and

(III), manganese (I) and (II), copper (I) and (II) and corresponding pairs of

ions from cerium, vanadium and iron. It has been suggested that the metal

ion may be interacting with the aldehyde to form an initiating radical accord-

ing to the equation

M+3 + RCHO M+2 + H+ + RCO- (5)

or with the peracid formed in the system via the following reactions:

M+2 + RCOOOH M(OH)+2 + RCO" (6)

M(OH)+2 + RCOOOH -- M+ + HOH + RCOOO (7)

Hence, if a metal in the lower oxidation state is added to a solution containing

peracid, the metal ion will be oxidized, then reduced, then oxidized again, etc.

Each of these reactions will generate a free radical that is capable of removing

a hydrogen atom from a molecule of aldehyde thus initiating the chain. Analo-

gous steps are postulated for the autoxidation of hydrocarbons.

It is reported that mixtures of ions of different metals are more effi-

cient catalysts than the individual ions taken separately (1).

The aim of this work is to study the autoxidation of benzaldehyde in an

effort to elucidate in more detail the mechanism of metal ion catalysis. Hav-

ing done this it might be possible to extend the work to explain the synergism

between two different metal ions in an autoxidation.

This work is limited to an investigation of the effects of manganese,

cobalt and cerium salts in the autoxidation of benzaldehyde in glacial acetic



The information reported herein was obtained by 1.) non-kinetics

studies which include identification of the reaction products and the isola-

tion and characterization of intermediates, and 2.) kinetics studies on the

rate of oxygen uptake, the rate of substrate loss and the rate of formation

and loss of the intermediate peracid as various parameters are altered.

Further kinetics studies were performed to establish the nature of the in-

termediate stages of the reaction in the absence of benzaldehyde using

manganese ions as a catalyst.

II. Autoxidation of Benzaldehyde

In 1832 Liebig (2) observed that benzaldehyde was converted to benzoic

acid on standing in the air and that the reaction was accelerated by light.

Other workers (3, 4) found that autoxidizing benzaldehyde was a strong oxidiz-

ing agent and that, when acetic anhydride was used as a solvent, twice as much

oxygen was consumed (4), yielding acetic acid and benzoylacetyl peroxide as

products. The rate of oxidation is decreased by an extensive number of inhi-

bitors and is increased by radical initiators and by traces of metals. Blckstrim

(5, 6) postulated that a chain mechanism was involved in the autoxidation of both

heptaldehyde and benzaldehyde after discovering quantum yields from 10,000 to

15, 000 for the reactions. Christianson (7) had previously suggested a chain

mechanism for these reactions because of the observed inhibitions.

Subsequently, this reaction and others have been studied by a large num-

ber of workers so that some detailed mechanisms of aldehyde autoxidation have

been established.

A. Non-kinetics studies

The first product of the benzaldehyde autoxidation is the peracid which

reacts heterolytically with unreacted aldehyde to give the final acid product in-

dicated by the equation



The rate of this reaction is slow (8, 9) so that in an autoxidation there will be

an increase in the concentration of the peracid as the reaction proceeds (10,

11, 12).

A complex, stable at -300, was isolated using peracetic acid and

acetaldehyde (13). It decomposed to acetic acid upon heating with manganese

(II) ion and was differentiated from peracetic acid by using a 10 per cent

potassium iodide solution. Iodine was liberated more rapidly by the peracid

than by the complex. Bawn and Williamson (14) proposed that a complex be-

tween two molecules of aldehyde and one molecule of peracid was also formed

in acetaldehyde autoxidation. Evidence from their studies indicated that the

decomposition of the complex gave one molecule of acetic acid and one mole-

cule of aldehyde.

A complex of this nature has not been isolated from benzaldehyde-

perbenzoic acid mixtures (15, 16), although the kinetics data of Wittig and

Pieper (16) suggested that such a complex formed rapidly and decomposed

slowly at room temperature. Two possible structures for the complex were

O0 0 OH
CgH5C .CHC6H5 and C6IH5COOC6 H5


*Two titrations were used to find the concentration of the complex; the
first, to find the peracid concentration and the second, to find the peracid and
complex concentrations.


A structure analogous to compound (II) has been proposed as an intermediate

in the Baeyer Villiger reaction by Criegee (17) and supported by Doering and

Dorfman (18).

Other products observed under various reaction conditions are carbon

dioxide (13) and a small amount of uncharacterized yellow solid (10, 19). With

inhibitors added, a variety of products were found, presumably resulting from

radical attack on the inhibitor (20, 21).

B. Kinetics studies

A variety of kinetics studies have been reported. The results of some

of these are summarized in Table 1. (22, 23) and in the following paragraphs.

Order in aldehyde. The dependence of the rate on the benzaldehyde

concentration was found to be second order in photoxidations (24), second order

in thermal oxidations at high oxygen pressure, first order in benzoyl peroxide

initiated reactions (12) and three-halves order in the presence of cobalt ions


Order in oxygen. Almquist and Branch (25) in 1932 found that the ther-

mal oxidation is first order in oxygen, but for the same reaction in benzene

solution there was not a simple oxygen dependency. However, as the benzalde-

hyde concentration was increased the reaction approached first order in oxygen


The order changes with both temperature and pressure. At 160 the

thermal oxidation is zero order, but as the temperature is increased a first-

order relation is established (27).



Initiation Oxygen order Experimental Reference
activation energy

Photo Chem. 0 1.8 27

Thermal 0 7.6 27

Thermal 0 13.6 12

Thermal 0-1 17.7 26

Photochemical 5. 24

Cobalt 0 14.7 8

Benzoyl peroxide 0 17.2 12


The oxidation rate was found to comply with the generalized expression


Rate = a ) 9)

which gives orders between zero and unity depending on the values of the para-

meters a, b, and c.

The cobalt ion-catalyzed oxidation was found to be zero order in oxygen

from 550 to 950 mm. pressure (8). The photoWxidation in decane solution has

also been reported to be zero order in oxygen.

Activation energy. Studies of the activation energy have been compli-

cated by the change of order in oxygen with temperature. Both the activation

energy of the benzaldehyde autoxidation and the effect of temperature on the

oxygen concentration in solution are reflected in the rate-temperature relation-

ship. If the oxygen order is shown to be zero, then the observed activation

energy from an Arrhenius plot will be equal to the real activation energy.

However, if the rate is dependent on the change of oxygen solubility with

temperature the apparent activation energy will consist of a chemical rate term

and an equilibrium term so that a correction must be made to obtain the activa-

tion energy from an Arrhenius plot. Table 1 summarizes the activation energies

reported by various workers.

Initiation. The autoxidation of benzaldehyde has been initiated in many

ways including radical sources, ozone, metal catalysts and photolysis. For

the thermally induced reaction there is considerable disagreement as to the

possible initiating step and some authors feel that the "uncatalyzed" reaction is

in fact initiated by traces of metal ions or peroxides. The metal ion catalyzed
initiation will be discussed in more detail.


Chain length. In a radical chain reaction, under steady-state conditions,

the number of chains initiated must equal the number of chains terminated.

BEckstrim and Beatty(19), using anthracene as an inhibitor in benzalde-

hyde autoxidations were able to get some measure of the chain lengths by assum-

ing that each molecule of anthracene oxidized during the reaction corresponded

to one chain termination. The number of molecules of benzaldehyde oxidized as

measured by the total oxygen consumption divided by the number of chains

terminated gave the average chain length. The chain lengths calculated are

very large. This may be accounted for by the fact that anthracene is a rather

ineffective inhibitor and other termination processes are occurring simultane-

ously. The results, however inaccurate, undoubtedly show that the autoxidation

is of chain character.

Briner and Papazian (28) determined the chain length for ozone-initiated

benzaldehyde oxidations in carbon tetrachloride solution to be as high as 5,000

using air and as high as 50,000 using pure oxygen.

Inhibition. Bafckstrom (5) showed that in the photo-initiated autoxidation

of benzaldehyde the order in light intensity depended upon the kind of inhibitor

used. Using anthracene as an inhibitor, the order in light intensity was found

to be 0.65, while for others the following values were obtained: benzyl alcohol,

0.5; diphenylamine and hydroquinone, 0.9. Pure benzaldehyde gives an order of

0.5 in light intensity.

Waters (29) studied the retardation of benzoyl peroxide-initiated benzalde-

hyde autoxidation. His kinetics results indicate thatp-cresol and m-2-xylenol

act primarily as chain-transfer agents and in view of the observed dependencies


of rate on the concentrations of benzaldehyde, benzoyl peroxide and inhibitor,

it was concluded that the main chain ending process was a dimerization, or

possibly a disproportionation, of the phenolic inhibitor radicals:

2 C7H70* -- inert products (10)

With the m-2-xylenol system, it was established that the dimer (C8H9O)2 is

further oxidized to 3, 5, 3', 5'-tetramethyl-4,4'-diphenoquinone. The di-

quinone is still a retarder of the autoxidation and it was inferred from further

kinetics studies that it acts as a chaining stopping agent by combining with

benzoyl radicals.

Later kinetics studies by Waters (20, 21, 30) of the benzoyl peroxide

catalyzed autoxidation, retarded by a number of polycyclic hydrocarbons, in-

dicated that chain termination is initially effected by combination of benzoyl-

peroxy radicals with the hydrocarbon.

The peracid-aldehyde reaction. The reaction by which the peracid goes

to the acid has been found to be first order with respect to both the peracid and

aldehyde (13) in the case of acetaldehyde. It was observed that the intermedia-

tory peracetic acid reacts with acetaldehyde to yield a peroxide compound

which was decomposed by manganese (MII) catalyst to give two molecules of

acetic acid. Bawn (15) has indicated that no complex is formed between per-

benzoic acid and benzaldehyde. However, there are some kinetics data that

suggest the rapid formation of a complex which decomposes slowly to benzoic

acid (16).

There has been more extensive work in the peracid oxidation of ketones

than of aldehydes. It is now generally accepted that the reaction is ionic in


character and proceeds by the following steps:

0 O O-H
R-C-R' + R-C-OOH - R-C-R'

0 0
R-0-C-R+ R-C-OH (11)

This scheme accounts for 1.) the isolation of hydroxyhydroperoxides which can

be decomposed by heating, 2.) the fact that the migratory aptitude of R and R'

is promoted by their electron releasing ability, 3.) the observation that in the

oxidation of acetophenone with perbenzoic acid the rate determining step is the

acid catalyzed addition of the peracid to the carbonyl group (31) and 4.) the

radioactive isotope data (18).

Benzaldehyde reacts with hydrogen peroxide in ether to give the correspond-

ing acid and 0.7 per cent phenol (32). Benzaldehydes substituted with electron

pushing groups when reacted with hydrogen peroxide under the same conditions

give higher yields of phenols. For example, 2, 4-dimethoxybenzaldehyde gives

26 per cent phenolic products and 74 per cent substituted benzoic acid. Phenolic

products produced as a side reaction could be responsible for the self-inhibition

of an autoxidation.

C. The reaction mechanism

A mechanistic rationalization of the kinetics and non-kinetics data for

aldehyde autoxidation has been formulated as more and more facts were deter-

mined. The first substantial break-through was showing that the reaction

proceeded via a free radical mechanism. Then the question as to the particular

nature of the individual steps became the subject of considerable effort and, at

times, considerable conjecture.


If the individual steps are to be resolved from idnetics data, it is

necessary to know the concentrations of the various radicals involved.

Until the present time, no method has proven powerful enough to measure

the very small radical concentration in an autoxidation with sufficient

accuracy to be helpful. Electron paramagnetic resonance has the calculated

theoretical sensitivity to detect 2 x 10"14 moles of free radicals (338 34), but

it has been largely limited to radical species that are stable in solution or

captured in the solid state. However, spectra have been observed of free

radicals involved in vinyl polymerizations (35) and in highly inhibited autoxi-

dations (36).

If it is assumed that the radical concentration is constant, i.e., a

steady state has been reached in which the rate of radical production equals the

rate of disappearance, it is then possible to write differential equations for the

rate of radical production and disappearance and set them equal to zero. In

this manner, it is possible to determine the ratios of the rate constants of

the separate reactions in the total sequence and to predict the order in sub-

strate and catalyst. If the result agrees with the experimental rate law, then

the mechanism proposed is a possible one. There is, of course, the chance

that the results coincide fortuitously. Thus, this type of agreement is neces-

sary but not sufficient for a proof of mechanism.

The initiation reaction. Photochemically induced reactions (37) are

primarily initiated by homolytic cleavage between the alkyl or aryl group and

the carbonyl group:




Cleavage between the carbonyl and hydrogen probably does not occur, since

attempts to isolate acetyl iodide from the products of acetaldehyde-iodide

photolysis have been unsuccessful (38). This is reasonable in light of the

higher bond dissociation energy of the C-H bond as compared with the C-C

bond (39).

Mulcahy and Watt (26) have studied the kinetics of uncatalyzed autoxi-

dation of benzaldehyde in benzene solution. The initial rate of oxidation is

given by the relation:

d O2 ] FRH]2 [O2 [
dt b[RHj +c[ 02] (1

The reciprocals of the constants b and c vary with temperature according to

the Arrhenius equation. The authors were able to conclude that:

1.) The mechanism of the thermally initiated

reaction is different from the various

catalyzed, photochemical and photosensi-

tized oxidation previously studied (14, 15,

40, 41).

2.) More work on the kinetics of uncatalyzed

thermal oxidations is desirable.

They proposed the following mechanism:

RH Q* initiation (14)

Q* + 0--- Q*02 (15)
Q*o + RH --- ROOH + Q* (16)

Q2 -- x (17)

Q + RH- -X termination (18)

where Q and QO2* are unspecified chain carriers. This mechanism is similar

to that suggested by George and Robertson (42) to explain the kinetics of tetralin

oxidation at 1100.

The major paths by which metal salts catalyze autoxidation appear to be

reaction with the peroxide intermediate (43, 44, 45, 46) and/or direct initiation

involving the aldehyde. With acetaldehyde the autoxidation rate is proportional

to aldehyde and cobalt concentrations but independent of the oxygen pressure (14).

Thus, the kinetics suggest:

Co+2 + CHI3CO3H Co+3 + CH3CO2 + OH- (19)

Co+3 + CHI3COt3 -H Co+2 + CH3CO3' + H+ (20)

as primary processes.

For benzaldehyde in acetic acid Bawn found the rate law to be:

d [02] k [RCHOj 3/2 [Con] 1/2 (
~ (21)

This rate law may be rationalized in terms of a direct initiation reaction.

Initiation of this type was first proposed by Haber and Willstttter (47)

for the iron (II) ion catalyzed autoxidation of acetaldehyde:

CH3CHO + Fe3+-- CH3CO' + H+ + Fe2+ (22)

Uri (48, 49) has suggested a third type of initiation in the cobalt ion catalyzed

oxidation of methyl stearate in which the cobalt ion acts as an oxygen carrier

capable of initiating chains:

Co+2 + 02 CoO2* (23)

CoO2.+2 + Co+2(RHI) --- Co+2 + Co+3R- + HO2* (24)

HO2. + RH -- HO2H + R (25)

Able (50, 51) suggested the same scheme for the lead catalyzed oxidation of


The propagation reaction. It is generally agreed that the chains are

propagated by:

RCO- + 02 RC03. (26)

RC03' + RCHO ----RCO3H + RCO. (27)

The termination reaction. There are three possible termination reac-

tions in uninhibited aldehyde autoxidation:

2RCO ---RCOCOR (28)

RCO' + RCO-" RCOO2COR (29)

2RCO3 --- (RC3)2 (30)

In an oxidation that is zero order in oxygen it is reasonable that the concentra-

tion of RCOg3 is large in respect to the last reaction.

Inhibitors may be added to the reaction mixture to react with one of the

chain-carrying radicals to shorten the average chain length thus decreasing the

reaction rate. There are two possibilities with phenolic inhibitors:

RCO- + R'OH RCHO + R'O' (31)

RCO3. + R'OH ---RCO3H + R'O (32)


If R'O" is consumed by a radical-radical reaction, the order in the initiating

reaction will be one. Intermediate orders would result from a competition

between the various termination reactions.

There is some evidence for the self-inhibition of liquid-phase photolysis

initiated autoxidation of benzaldehyde (24, 52). It was suggested that a com-

pound corresponding to the formula


was responsible. In the cobalt catalyzed benzaldehyde oxidation with no solvent,

water and phenols formed during the reaction retarded the process (53). The

introduction of strong initiator decreased the retarding effect to a certain


Termination reactions that involve metal ions have been suggested (54,

55, 56) in the ferrous-ion-peroxide-initiated polymerization of acrylonitrile

and similar monomers. There appears to be competition between the termina-

tion reaction:

RO- + Fe2+-- RO" + Fe3+ (33)

and the propagation reaction:

RO' + CH2 =CHR' -- ROCH2-CHR (34)

Bamford et al., (57) have shown that ferric chloride is an efficient in-

hibitor of vinyl polymerization and Hammond (39) has found that ferric chloride

inhibits the autoxidation of tetralin in chlorobenzene.


Metal ions reacting with radicals can not be excluded as a possible

termination reaction in aldehyde oxidations (56, 58).


I. General Statement of Experimental Procedure

The rate of autoxidation of bonzaldehyde in glacial acetic acid under the

influence of metal-ion catalysts was studied in the apparatus described in the

following section. The oxidation rate was followed by observing the uptake of

oxygen as a function of time. From these data, the length of induction periods,

the rate of the reaction and the extent of the reaction were determined as func-

tions of various parameters.

In some runs the rate of loss of benzaldehyde was followed. In other

runs the rate of the formation of perbenzoic acid was found. The products of

the autoxidation were identified for selected runs.

II. Purification of Benzaldehyde

The general procedure for the purification of benzaldehyde for use in

kinetics experiments has been to perform repeated vacumn distillations under

an inert gas such as hydrogen or nitrogen (26, 59).

Benzaldehyde (Merck) was allowed to stand one week over anhydrous

potassium carbonate. Purified nitrogen was kept over the benzaldehyde to

prevent oxidation. The benzaldehyde was filtered and transferred to the dis-

tillation flask together with fresh anhydrous potassium carbonate.

The distillation was performed using a one meter column equipped

with a variable take-off head controlled by an electronic repeat-cycle timer.

The externally heated column was packed with single-turn glass helices. The

distillation flask was fitted with a nitrogen bleed capillary. A "fraction-

cutter" was also used so that the entire distillation could be run without

exposing the system to the atmosphere. After maintaining total reflux for

one hour 3,000 ml. of benzaldehyde was collected at a reflux ratio of six to

one. A 2,300 ml. center-cut was collected and redistilled at a reflux ratio of

ten to one. A center-cut of 1,800 ml. was collected from the second distilla-

tion and subjected to both infrared and gas phase chromatographic analysis.

The samples for infrared were run in 0.025 mm. cells on a Perkin-

Elmer Model 21 Infrared Spectrophotometer and there was no peak at 2.9


microns indicating that no appreciable amounts of water, benzoic acid or

perbenzoic acid were present.

The benzaldehyde was stored in a blackened pyrex flask fitted with a

delivery tube and a gas inlet. Removing benzaldohyde from the storage flash:

was accomplished by applying a positive nitrogen pressure through the gas


All glass apparatus was used in both the preparation and storage of

the benzaldehyde. Dow-Corning high vacuum grease was used on all ground

glass joints.

II. Preparation and Purification of Other Chemicals

A. Acetic acid

Baker reagent grade glacial acetic acid was distilled at atmospheric

pressure using the column and distillation head described under the purifica-

tion of benzaldehyde. Prior to the start of the distillation one per cent of

acetic anhydride was added to the distillation flask. After total reflux for two

hours 100 ml. of forerun was discarded, the center-cut was collected at a

reflux ratio of ten to one and 400 ml. was left in the distillation flask. The

system was kept under purified nitrogen throughout the distillation.

B. Manganese (II) acetate

Hydrated manganese (II) acetate, reagent grade, was obtained from

Fisher Scientific Company. A weighed amount of the compound was dried

over phosphorus (V) oxide at 1030 and 1.5 mm. of Hg for 72 hours. The

weight loss indicated that the product was anhydrous.

C. Manganese (III) acetate

Manganese (III) acetate was prepared by adding 80 g. of acetic anhydride

to 20 g. of 50 per cent manganese (II) nitrate in a magnetically stirred 2000 ml.,

three neck flask fitted with a reflux condenser, addition funnel and thermometer.

The manganese (II) nitrate solution was heated to 1450 before the anhydride was


added drop-wise. A temperature of 120 to 1250 was maintained without further

heating by regulating the addition rate of the anhydride, When the evolution of

nitrogen dioxide stopped the reaction mixture was allowed to cool to room tem-

perature and then it was placed in an ice bath for one hour. The solid manganese

(II) acetate was then filtered from the acetic acid and excess acetic anhydride.

The product was washed six times with acetic anhydride followed by eight wash-

ings with dry ether. The ether was removed from the light tan crystals under

reduced pressure at room temperature.

The equation for the overall reaction between manganese (II) nitrate and

acetic anhydride is:

2 Mn(NO3)2(H20)6 + 15 (CH3CO)20

2 Mn(CH3COO)3 + 4 NO2 + 1/2 02 + 24 CH3OOH (35)

The method used was modified from the procedure used by Chretien and

Varga (60).

D. Cobalt (II) naphthenate

Cobalt (II) naphthenate was obtained from the Noudex Corporation as

"Noudex Cobalt 6". It was found to contain 5.40 per cent cobalt as cobalt (II)

naphthenate in naphthenic acid (61). A solution 3.00 x 10-1 M cobalt in redistil-

led glacial acetic acid was prepared and analyzed spectrophotometrically. The

concentration of this solution was found to be 3.00 x 10-1 M cobalt.

*The writer is indebted to Mr. K. Hickey for the analysis of the cobalt

E. Cerium (IV) naphthenate

Cerium (IV) naphthenate was obtained from the Noudex Corporation as

"Noudex Cerium 6". Stock solutions were made up in glacial acetic acid.

F. Oxygen and nitrogen

Cylinder Oxygen (U. S. P.) was obtained from Linde Company. Passing

the oxygen through a long tube packed with anhydrous calcium chloride removed

possible traces of water.

Gas Phase Chromatography of the oxygen showed only one peak. One

cylinder was used for all of the autoxidation runs. No attempt was made to re-

move possible traces of ozone. However, no ozone was observed on the vapor


Prepurified nitrogen was obtained from Airco Company and passed

over hot copper to remove traces of oxygen before use as an inert atmosphere.

Oxygen was undetectable in the purified nitrogen by vapor phase chromatography.

G. Benzoic acid

Benzoic acid (U. S. P.) was obtained from Baker Chemical Company. It

was recrystalized three times from distilled water and dried for five hours at

room temperature under reduced pressure (m. p. 122.10).

H. Perbenzoic acid

Perbenzoic acid was prepared by the method of Braun (62) using the

modifications suggested by Kolthoff (63). An 85 per cent yield was obtained.

The white crystalline solid was dried for several hours under 10 mm. of Hg at

300 and stored at -5.


Prior to using the perbenzoic acid a sample was titrated to determine

the amount of active oxygen so that a correction could be made for decomposi-


I. Other chemicals

All other chemicals used were C. P. grade or better.

IV. Oxidation Apparatus

The oxidation apparatus used in this investigation is shown in Figure

1. The reaction flask was maintained at constant oxygen pressure by feeding

the gas from a supply tank when oxidation of the substrate occurred, A mer-

cury manometer controlled a relay that actuated a solenoid valve bleeding

oxygen from the supply tank to the reaction flask. The reaction flask pressure

could then be maintained to + O~ 5 mm, of Hg,

The pressure drop in the supply tank was used to measure the oxygen

uptake. This pressure was recorded by a Taylor "Fulscope" Recorder which

consists of a set of bellows that plot pressure as a function of time on a cir-

cular chart. The chart records the per cent of oxygen used,

The relationship between per cent oxygen and the actual moles of

oxygen was determined by two methods. One method consisted of determining

the volume of oxygen bled from the tank at various per cent readings by collect-

ing the gas over water. The second method utilized was to absorb oxygen from

:the tank into a solution of alkaline pyrogallol. The weight increase of the

pyrogallol solution was then determined.

The results of both determinations show that the pressure recorder is

linear and that one per cent on the recorder chart corresponds to 1.40 mM of

oxygen (Figure 2).

Figure 1. Oxidation apparatus


Figure 1. Oxidation Apparatus (continued)

List of Parts

1. Taylor "Fulscope" Recorder model 455M

2. Azco solenoic valve no. 6867E

3. Bellows type needle valve

4. Four mm. glass tubing

5. Oxygen inlet for flushing the system

6. Bleed valve relay

7. Vent valve

8. Thermometer and thermoregulator

9. Sample removal valve

10. Bath temperature relay

11. Stirrer pump

12. Calrod heater, 250 watt

13. Copper tubing

14. Inlet valve from oxygen cylinder

15. Pressure regulator manometer

16. Insulated constant temperature baths maintained within + 0.070

17. Oxygen supply tank

18. Pyrex reaction flask

19. Teflon magnetic stirrer

20. Waco stirrer motor, 300 rpm

0 by volume

* by weight

2 4 6 8 10 12
mM 02

Figure 2. Calibration of recorder








V. Typical Autoxidation Run

A. Sample preparation

In general, 25 ml. of benzaldehyde was added to 200 ml. of glacial

acetic acid containing the metal catalyst. The total volume was then brought

to 250 ml. by adding more acetic acid. The sample was transferred to the

reaction flask which was connected to the system and immersed in the constant

temperature bath with the sample vent open. The system was flushed with oxy-

gen for two minutes at 5.0 liters per minute, after which the recorder was

started. One minute after starting the recorder the sample vent was closed.

Covering the bath prevented light from reaching the reaction flask.

Solid additives were weighed out on a Mettler balance (sensitivity

0.02 mg.) and dissolved in 200 ml. of acetic acid. Additive solutions were

pipetted into 200 ml. of acetic acid prior to adding the substrate.

All equipment was thoroughly cleaned before use to avoid contamination

from preceding runs in the following way:

After allowing the reaction flask to stand in concen-

trated sodium hydroxide solution for 24 hours it was

rinsed with tap water. This was followed by rinsing

six times with distilled water and six times with re-

distilled acetone. Purified nitrogen was used to dry


the flask and other equipment. Pipettes were

cleaned with three charges of acetone, three

distilled water rinses and six rinses with re-

distilled acetone.

B. Sample removal

Withdrawing samples after the run had started to determine the concentra-

tion of benzaldehyde and benzoic acid was accomplished by inserting a pipette

down the sample vent into the reaction mixture. The pipette was designed so

that it was small enough to fit through the sample valve. Samples were removed

only when the system was under one atmosphere.

C. Interpretation of autoxidation data

The data from the recorder chart were transferred to rectangular coor-

dinates. For example, the results of a typical run with 0.988 M benzaldehyde

and 8.4 x 10-4 M cobalt (II) ion at 500 are shown in Figure 3. The initial rate

was determined from this curve by plotting log(-dO2/dt) as a function of log [BzH]

and extrapolating to the initial benzaldehyde concentration, where -dO2/dt is the

oxidation rate and [BzH] is the concentration of benzaldehyde.

In a number of cases induction periods were observed. The length of the

induction period was found by the method of intercepts (61). A plot of a run

0.988 M benzaldehyde and 4.0 x 10"4 M cerium (IV) ion at 500 is shown in

Figure 4. The induction period, ti, was found to be 77 minutes.



30 -



10 -

.-.. ........._.. .__ i_ I
15 20 25 30 35

time (min. )

Figure 3. Typical autoxidation run ( 0.988 M BzH, 8.4 x 10-4 M Co(II),
500, 1 atm. 02 )


" 20 -



20 40 60 80 100 120 140
time ( min. )

Figure 4. Determination of the induction period ( 0.988 M BzH, 4.0 x 10-4 Ce(IV), 500, 1 atm. 02 )

VI. Analytical Methods

A. Peracid determination

Perbenzoic acid was estimated by titrating the iodine liberated from

the reaction between the peracid and potassium iodide (95) using the follow-

ing method: A five ml. sample was run into 25 ml. of 1.0 N sulfuric acid

containing 2.0 g. of potassium iodide and the liberated iodine was titrated

after three minutes with 0.984 N thiosulphate solution. A starch indicator

was used to show the end point.

B. Benzaldehyde determination

The benzaldehyde concentration was found using a Perkin-Elmer gas

phase chromatograph. Samples of constant volume (4 microliters) were in-

troduced into the column with a ten microliter syringe equipped with a device

to insure the delivery of equal sample volumes (Chaney adaptor). The area

under the benzaldehyde peak was found to be proportional to the concentra-

tion of the aldehyde and was used as a measure of the concentration for the


C. Product identification

Soluble products. The solution from an autoxidation was filtered to

remove any solid precipitate. The solution was then subjected to gas phase


chromatography. Retention times were used to determine what compounds

had been eluted from the column. The peak heights were used to estimate

the concentration of the products.

The reaction mixture was also subjected to paper chromatographic

analysis. The excess acetic acid was stripped off at reduced pressure and

a small amount of the solid residue was chromatographed using butanol-

pyridine-water (10:2:1) as the eluting solvent. The compounds were visualiz-

ed on the chromatographic paper by spraying sections of the paper with one

per cent ferric chloride, starch-iodide and ammoniacal silver nitrate solu-


Insoluble products. A yellow solid which precipitated from the man-

ganese ion catalyzed autoxidation of benzaldehyde was subjected to extensive

physical and chemical tests.

The results of the product identification will be discussed in the Re-

sults Section.


I. Manganese Systems

The results in this section show the effects of benzaldehyde, manganese,

water and acetic anhydride concentrations and oxygen pressure on the autoxida-

tion. Both manganese (II) and (mI) ions were used to catalyze the reaction. The

dependency of the rate on temperature was determined. Experiments were also

performed to determine the effect of the oxidation products on the reaction.

The metal ions were added by introducing a known volume of stock

catalyst solution in acetic acid or in some cases weighing out the salt as a solid.

The rate of oxygen uptake was determined in the oxidation apparatus and the

initial rate and the length of the induction period were then found by the method

described in the Experimental Section. Unless otherwise stated the reaction

temperature was 500.

Several general observations concerning the course of the reaction can

be made:

1. The color of the metal ions was undiscernible in the concentrations

used. However, just prior to the uptake of oxygen the solution turned brown.

This effect was most pronounced in the runs with induction periods. The pro-

duction of the brown color invariably preceded the start of the oxidation. The

brown color results when perbenzoic acid and manganese ions are in the system.


2. Oxygen uptake occurs smoothly in the presence of manganese ions

but stops abruptly at a time dependent upon the manganese ion concentration

and certain other factors.

3. At low concentrations manganese ions act as catalysts. However,

at higher concentrations this catalytic effect disappears and a retarding effect

is observed,

A. The effect of manganese (I) ion concentration

A series of autoxidation runs at constant benzaldehyde concentration

and total volume was performed to establish the effect of manganese (II) ion on

the reaction (Table 2). An increase in the oxidation rate was observed as the

manganese ion concentration was increased from 2 x 10-5 M to about 2 x 10"4

M. A further increase in the manganese ion concentration resulted in a de-

creased autoxidation rate. These effects are shown in Figure 5.

An induction period, proportional to the manganese ion concentration,

was observed at manganese concentrations greater than 5 x 10"4 M (Figure 6).

The uptake of oxygen was observed to stop before all of the benzalde-

hyde was consumed. The limited uptake of oxygen, called the extent of

reaction, is expressed quantitatively as the number of moles of oxygen consumed

in the reaction. The extent of the reaction increased and then fell off as the man-

ganese ion concentration was increased as shown in Figure 7.

B. The effect of manganese (II) ion at low concentrations

From the data obtained using a wide spread of catalyst concentrations,

it is obvious that the effect of manganese ion is not a simple process, but possi-

bly a combination of both catalysis and inhibition.



( [BzH] o= 0.988, 500o 1 atm. 02)

[Mn (II)] d02/dta Extent of reaction ti

Mx 105 mM 02/min. mM 02 min.














































a initial post induction oxidation rate


0.5 -

without acetic anhydrid
S0.4 -

without acetic anhydride


0.2 -

2 4 6 8 10

[Mn(II)] x 104

Figure 5. The effect of manganese (II) ion on the oxidation rate of
benzaldehyde ( 0.988 MBzH, 500, 1 atm. 02 )





4 O



0.6 1.0 1.5 2.0

[Mn(II)] x 103

Figure 6. Induction periods for manganese catalyzed autoxidation of
benzaldehyde ( 0.988 MBzH, 500, 1 atm. 02 )




2 4 6 8 10

[Mn (II)] x 10

Figure 7. The effect of manganese (II) ion on the extent of reaction
( 0.988 MBzH, 500, 1 atm. 02 )


At manganese ion concentrations of less than 2 x 10-4 M the metal

ion catalyzes the autoxidation, A plot of the autoxidation rate against the

square root of the catalyst concentration was linear from the lowest man-

ganese (II) ion concentration, 2 x 10-5 M, to 2 x 104 M manganese ion, as

shown in Figure 8. This indicates that the manganese-initiated reaction is

one-half order in metal ion,

C. The effect of manganese (II) ion concentration

A series of experiments was performed using manganese (I) ions

rather than manganese (II) ions as a catalyst. The dependency on manganese

(I) was found by holding the concentration of benzaldehyde and the total vol-

ume constant while the catalyst concentration was varied over the range,

8 x 10-6 to 4 x 10-3 M. The results are summarized in Table 3.

In general, the autoxidation rate exhibited a similar dependency on

manganese (III) ions as it did on manganese (II). At low concentrations of

the metal ion (less than 4 x 10-4 M) the rate increased as the catalyst concen-

tration was increased, but it then leveled off and finally, at higher manganese

(III) concentrations became inversely proportional to the concentration.

The autoxidation rate using manganese (II) ion was significantly

greater, under the same conditions, then the rate using manganese (II) as a

catalyst. Thus, manganese (I) appears to be more effective in initiating the

autoxidation of benzaldehyde than manganese (II).

The effect of manganese (III) ion on the extent of reaction paralleled

that of the manganese (II) system. The extent of reaction at first increased

then decreased as the manganese (III) concentration was increased (Table 3).



0 0.4



1 2 3 4 5

[Mn (II)] 2x 105

Figure 8. The effect of manganese (II) ion concentration on the rate
of oxidation ( 0.988 MBzH, 500, 1 atm. 02 )



( [BzH] o=0.988, 500, 1 atm. 02)

[Mn(m)] -dO2/dta ti Extent of reaction

Mx 105 mM O2/min. min. mM 02

0.8 0.379 0 75.6

2.5 0.516 0 90.3

45.8 0.798 5 121.8

200 0.369 69 112

400 0.229 306 90

initial post induction rate


For equal catalyst concentrations, the extent of reaction was greater for man-

ganese (II) than for manganese (II).

The contrast in the extent of reaction between the two catalysts may be a

ramification of the relative efficiencies of initiation and/or inhibition by man-

ganese (II) and (m) ions.

Induction periods were observed at high manganese (I), ion concen-

trations (Table 3). They are somewhat shorter than the induction periods

observed in the manganese (I) systems.

In summary, it has been shown that:

1. Manganese ions display both inhibition and initiation of the autoxida-

tion of benzaldehyde with the latter predominating at low manganese ion

concentrations and both processes occurring as the metal ion concentration


2. At high metal ion concentrations the length of the induction period is

proportional to the manganese ion concentration, with manganese (TI) giving

longer induction periods than manganese (III).

3. Both manganese (II) and (III) behave in a similar manner, but man-

ganese (III) is the more effective catalyst.

D. The effect of perbenzoic acid

The dependency of the manganese ion catalyzed autoxidation of benzalde-

hyde on the concentration of perbenzoic acid was studied by adding known amounts

of the solid peracid at the start of the run and following the rate of oxygen uptake.

The results are shown in Table 4 and indicate that for low manganese ion con-




(PzH% = 0.988, [Mn(n)]= 4.0 x 10-5, 500, 1 atm. 02)

BzOOHi -dO2/dt

Mx 104 mMOl2/min.

1.0 0.363

2.5 0.387

5.0 0.424

10.0 0.354

50.0 0.338

100 0.296

500 0.373


centrations (4.0 x 10-5 M) small amounts (less than 5 x 10-4 M BzOOH) of

perbenzoic acid will initiate the oxidation. However, this initiation reaches

a maximum at about 5 x 104 M perbenzoic acid and further increases in the

peracid concentration result in a decrease in the reaction rate.

A plot of -dO2/dt versus [BzOOH] 2 is linear over the concentra-

tion range between 1 x 104 and 5 x 10-4 M (Figure 9). This indicates that

there is a one-half order dependency on the initial concentration of perbenzoic


The following experiments were performed to examine the effect of

both high metal ion and perbenzoic acid concentrations. A reaction mixture

containing 2.0 x 10-3 M manganese (II) and 1.2 x 10-2 M perbenzoic acid had

an oxidation rate of 0.13 mM 02/min. and no induction period. An equivalent

run without peracid had an oxidation rate of 0.31 mM O2/min. and a 73 min.

induction period. This indicates that an initially high peracid concentration

combined with a high metal ion concentration results in a severe retardation

of the oxidation. A peracid determination on the above reaction shows that the

added peracid was quickly decomposed and that there was no appreciable build-

up of peracid during the reaction.

E. The effect of benzaldehyde concentration

The rate of oxygen uptake was followed for reaction mixtures varying

in benzaldehyde concentration, but all at 500, and containing 2.0 x 10-5 M

manganese (II) acetate.



* 0.400



I - I
1.0 1.5 2.0 2.5

[BzOOH 'x 102

Figure 9. Plot showing the dependency on perbenzoic acid
( 0.988 MBzH, 500, 1 atm. 02 )


The oxidation rate increased with the benzaldehyde concentration as

shown in Table 5 and Figure 10. A plot of log(-dO2/dt) versus log [BzH]

(Figure 11) was not linear suggesting that the order of the reaction in ben-

zaldehyde changes with concentration, being about 1.3 order at low benzalde-

hyde concentrations and 0.2 at higher concentrations.

F. The effect of oxygen pressure

The dependency of the autoxidation rate on the oxygen pressure was

examined by carrying out a series of runs at constant aldehyde and catalyst

concentrations at 500.

The results of these experiments show that within experimental error

the autoxidation rate was independent of oxygen pressure (Table 6).

G. The effect of water

A series of experiments to determine how added water would alter the

reaction was performed. Both the benzaldehyde and manganese (II) ion con-

centrations were held constant while various amounts of water were added to

the reaction mixture,

The autoxidation.rate was found to be inversely proportional to the

water concentration (Table 7). The rate falls off sharply as small amounts of

water are added; this effect then levels off so that a continued increase in the

water concentration results in only small changes in the rate. However, add-

ing water increases the extent of the reaction (Figure 12). Although no

quantitative determination was made it was observed that the amount of yellow

solid formed during the reaction decreased with increasing concentrations of



( [Mn(I)] = 2.0 x 105, 500, 1 atm. 02)

[BzH] o -dO2/dt mM 02/min.

0.198 0.310a

0.394 0.168

0.593 0.322

0.789 0.370

0.988 0.472

rate after an induction period of 366 min.




-0.3 -

0. 2

0.2 -


2 4 6 8 10

[BzHil x 101

Figure 10. The effect of initial benzaldehyde concentration on the
oxidation rate ( 2.0 x 10-5 M Mn(II), 500, 1 atm. 02 )



4 0.4


0.3 -


0.6 0.7 0.8 0.9 1.0

log ( [BzH] x 101 )

Figure 11. Plot of log dO2/dt versus log [BzH]



( [BzH] =0.988, [Mn(II)] =4.0ox105, 500, 1 atm. 02)

Total pressure






mM 02/min.





___ I~



( [BzH] =0.988, [Mn(II)] =4.0x10-5, 500, 1 atm. 02)

H2 ] -dO2/dt Extent of reaction

M x103 mM O/mln. mM 02
- -- ^' -- 2



















a [Mn(I)] =2.0 x 10-5



80 --

70 -




2.5 5.0 7.5 10.0

[H20l x 102

Figure 12. The effect of water on the extent of reaction ( 0. 988 M BzH,
4. 0 x 10-5 M Mn(II), 500, 1 atm. 02 )


water. No yellow solid was formed when the reaction mixture was 2.0 x 10-5 M

manganese (II) ion and 1.11 x 10-1 M water.

H. The effect of oxidation products

The effect of solid product. The solid product from an autoxidation run

was washed six times with glacial acetic acid and added to a run containing 0.988

M manganese (II) acetate. Both the rate of oxidation and the extent of reaction

were identical with that of a run containing equal substrate and catalyst concen-


The effect of soluble products. A run containing 0.988 M benzaldehyde

and 2.0 x 10-5 M manganese (I) ion, was allowed to proceed for 24 hours. At

this time a 10 ml. sample from this run was withdrawn and filtered and added to

a second run of equal substrate and catalyst concentration. The second run ex-

hibited a long induction period after which some oxygen was slowly absorbed by

the system. The maximum rate of oxidation for this run was 1.1 x 10-2 mM

O2/min. as compared with a rate of 3.5 x 10-1 mM O2/min. for a run of identi-

cal benzaldehyde and catalyst concentration.

The effect of benzoic acid. The effect of benzoic acid on the manganese

(II) ion catalyzed reaction was measured by adding known concentrations of the

acid to the reaction mixture. The results are shown in Table 8. These data

indicate that the oxidation rate was unchanged for different initial benzoic acid


I. The effect of catalyst on the peracid concentration

The concentration of perbenzoic acid was followed iodometrically during

the autoxidation of benzaldehyde at a number of catalyst concentrations. The re-



( [BzH ]o 0.988, [Mn(II) =4.0 x 105,. 500, 1 atm. 02)

[BzOH] 0

M x 102


mM 0 mmn.

0.0 0.372

1.0 0.372

5.0 0.376

_ __
I _


suits as shown in Figure 13, indicate that the peracid level increases steadily,

reaching a maximum, after which it slowly decreases. As the catalyst concen-

tration was increased the maximum concentration of perbenzoic acid obtained

during the run was lowered.

J. Peracid concentration during induction period

The build-up of perbenzoic acid was followed for a run that exhibited an

induction period due to the high concentration of manganese (II) ion. The reac-

tion mixture was 0.988 M benzaldehyde and 2.0 x 10-3 M manganese (II). The

peracid level was observed to increase slowly during the initial part of the ini-

tiation period, followed by a sharp increase at the start of rapid oxygen uptake

and then a slow decrease in concentration (Figure 14). The perbenzoic acid

build-up during the induction period appears to be exponential as shown in

Figure 15.

K. The effect of temperature

Experiments carried out at different temperatures with reaction mixtures

containing 0.988 M benzaldehyde and 2.0 x 10-5 M manganese (I), show how the

autoxidation rate was changed with temperature (Table 9).

The Arrhenius plot of these data was found to be non-linear as shown in

Figure 16. Non-linearity can result if there are two competing reactions of

different activation energies occurring. The rate fell off at higher temperatures

suggesting that the inhibition reaction has a higher overall activation energy than

the initiation reaction.

[Mn (I)]

0- 0.0

4 Q 2.0 x 10-5

0 4.0 x 0-5

@ 2.0 x 10-3


0 0


_ r It

100 200 300 400 500 900 1000

time (min. )

Figure 13. Peracid concentration during the autoxidation at various
catalyst concentrations (0.988 M BzH, 500, 1 atm. 02 )

3 ti

0 /

40 80 120 160 200 240 280 320

time (min.)

Figure 14. Increase in peracid concentration during the induction
period ( 0.988 MBzH, 2.0 x 10-3 M Mn(II), 500, 1 atm. 02)






o. 6 0.



20 40 60 80

time (min. )

Figure 15. Exponential build-up of peracid during the induction
period ( 0.988 MBzH, 2.0 x 10-3 M Mn(II), 50o, 1 atm. 02 )



( [BzH] =0.988, [Mn(II)] =2.0 x 105 1 atm. 02)

Temperature dO/dt mM O2/min.

30.01 0.170

34.50 0.295

40.00 0.371

50.00 0.410

60.00 0.420






a 0.3




3.0 3.1 3.2 3.3 3.4

1/T x 103

Figure 16. Arrhenius plot for manganese-catalyzed benzaldehyde
autoxidation ( 0.988 M BzH, 2.0 x 10-5 M Mn(II), 1 atm. 02 )

L. The effect of acetic anhydride

If acetic anhydride is added to the autoxidation run the effect of chang-

ing the manganese (II) ion concentration is altered as shown in Table 10. In

this experiment 1.0 ml. of redistilled acetic anhydride was added to the reac-

tion mixture before the start of the run and the rate of oxygen uptake was

followed at 500.

The effect of the change in the metal ion concentration on the oxidation

rate with and without added acetic anhydride is shown in Figure 5. The results

indicate that added acetic anhydride had little effect on rate of oxygen uptake

when the manganese (II) ion concentration was less than 2.0 x 10"4 M. How-

ever, at higher catalyst concentrations, the oxidation rate for runs with added

acetic anhydride was considerably higher than the rate without the anhydride.

M. Loss in benzaldehyde

The course of the reaction was followed by the loss in benzaldehyde, the

oxygen absorbed and the peracid formed. The analytical data for a typical run

are shown in Table 11.

The results show that oxygen uptake and loss of benzaldehyde stopped at

the same time; however, the concentration of perbenzoic acid was at a maximum

at the time the uptake of oxygen stopped, after which it slowly decreased as shown

in FIgure 17.

A change in the perbenzoic acid concentration resulting from a reaction

with benzaldehyde would not be reflected by a like change in the aldehyde concen-

tration, due to the insensitivity of the gas phase chromatographic determination

of the benzaldehyde.



( [BzH] =0.988, 500, 1 atm. o.)

[Mn(n)] dO2/dt Extent of reaction

Mx 105 mM O0/min. mM O

0.0b 0.270

0.0 0.270

4.0 0.349 44.6

4.0 0.349 51.6

8.0 0.419 99.0

20 0.509 139.5

40 0.465 120.0

100 c 0.409

al.0 ml. of acetic anhydride per 250 ml. of solution
no acetic anhydride added

Induction period of 6.0 min.



( [BzH] =0.988, [Mn(II)] =4.0x10-5, 500, 1 atm. 02)

Reaction time [BzH ] [BzOCH] 02 absorbed dO2/dt

minutes M M x 103 mM mM 02/min.

0 0.99 0.0 0.0 0.372

65 0.81 2.36 21.3 0.280

120 0.70 4.83 36.8 0.245

175 0.63 10.00 51.8 0.230

55 0.59 6.15 58.8 0.0

300 0.59 3.38 58.8 0.0

360 0.59 1.78 58.8 0.0

oxygen absorbed

perbenxoic acid


60 120 180 240 300 360

time ( min. )

Figure 17.
Mn(II), 500,

1 atm. 02 )

of benzaldehyde ( 0.988 MBzH, 4.0 x 10-5 M


The same general results were obtained over the entire range of

catalyst concentrations (2 x 10-5 to 2 x 10-3 M Mn(I) ).

N. Some observations concerning the extent of reaction

Since the autoxidation stops before all of the benzaldehyde is consumed

the following experiments were performed to establish the factors involved in

the retardation:

1. The addition of fresh manganese (I) ion had no effect on the reac-

tion after it had stopped.

2. The addition of benzoyl peroxide and 2, 2'-azobisisobutyronitrile to

a reaction that contained unreacted benzaldehyde promoted a small uptake of


3. The addition of cobalt (II) naphthenate had a pronounced effect.

When the solution was made 1.2 x 10-4 M in cobalt a rapid uptake of oxygen was

observed after a 30 minute induction period. The oxidation rate after the addi-

tion of the cobalt solution was 0.20 miM 02/min. The benzaldehyde concentra-

tion at the point the cobalt (II) naphthenate was added was found to be 0.38 M.

The autoxidation rate of an independent run 1.2 x 10"4 M in cobalt at a benzalde-

hyde concentration of 0.38 M was found to be 0,256 mM 02/min.

The existence of the induction period indicates that when the cobalt (I)

naphthenate was added radicals were being produced but that an inhibitor was

preventing the initiation of the chain sequence. However, after exhausting the

inhibitor from the system the radicals introduced by the cobalt catalyst were

able to propagate the reaction chain at a rate nearly equal to that of an identical

reaction mixture without the inhibitor.


0. Product analysis of manganese-catalyzed benzaldehyde autoxidation

A number of runs were analyzed to determine the products of the reac-

tion. The products found include benzoic acid, perbenzoic acid, a trace of

water and a yellow solid.

Gas phase chromotography of the runs invariably gave four peaks. The

peaks were identified as oxygen, water, acetic acid and unreacted benzaldehyde.

Neither benzoic acid nor perbenzoic acid were eluted from the column.

The gas phase chromatographic analysis showed that water was pro-

duced in trace amounts which increase as the reaction proceeds. The final

water concentration estimated from peak areas was about 1 x 10-3 M for a run

containing 0.988 M benzaldehyde and 2.0 x 10-5 Mmangaganese ion and 5 x 10"4

M for a run 0.988 M benzaldehyde and 4.0 x 10-5 M manganese ion. Both runs

were at 500 and 1 atm. oxygen.

Paper chromatographic separations using pyridine-butanol-water as a

solvent gave three products, unreacted benzaldehyde (visualized with Tollen's

reagent, Rf = 0.13), benzoic acid (Rf = 0.96) and a trace of a phenolic or enolic

product (visualized with one per cent ferric chloride solution, Rf = 0.81).

The insoluble product from the manganese ion catalyzed autoxidation of

benzaldehyde had a melting point greater than 3500. Elemental analysis* showed

the compound to contain 70.66 per cent carbon, 3.96 per cent hydrogen and 25.38

per cent oxygen (calculated for CyH502: C = 69.42, H = 4.16 and O = 26.42 per


Analysis were performed by Galbraith Microanalytical Laboratories,
Knoxville, Tennessee.


No ash was obtained on heating. The solid was soluble in 5.0 per cent sodium

carbonate solution and in dimethylformamide and insoluble in acid and in cer-

tain polar and non-polar solvents.

The infrared spectra of the substance showed absorption at 1025, 1060,

1200, 1350, 1450, 1590, 1680, 1730, 3070 and 3460 cm-1.

When the solid was reacted with hot 50 per cent nitric acid benzoic acid

was isolated from the solution. Reaction with methyl magnesium bromide gave

a gas and a second solid, which upon hydrolysis gave a product having no carbonyl

absorption band (1680 cm-1) in the infrared. The compound would not react with

potassium iodide under acid conditions, thus ruling out the possibility of it baing

a peroxide.

Since acetic acid appears to be relatively inert towards the attack of

benzoyloxy radicals (64) it is unlikely that the solid was produced by such attack

on the solvent. There is, however, much evidence that indicates radical attack

on the aromatic ring occurs with ease (65, 66, 67).

On the basis of these observations the following structure is suggested

for the solid product:


C6g-C-O-/ O-C-CgH-6


II. The Catalytic Decomposition of Perbenzoic Acid

The decomposition rate of perbenzoic acid catalyzed by manganese (II)

and (III) ions was studied in acetic acid in the absence of benzaldehyde. Stock

solutions in glacial acetic acid of perbenzoic acid and manganese (II) and (I)

acetate were used to make solutions of different concentrations for the kinetics


The rate of decomposition was followed by withdrawing 5. 0 ml. samples

at fixed time intervals and determining the peracid concentration iodometrically.

The samples were shaken in a 500 water bath. No attempt was made to exclude

oxygen from the system.

On adding either manganese (II) or (III) ions to the solution of perbenzoic

acid a light brown color formed with an intensity proportional to the metal ion

concentration. The same color remained during the entire reaction. A solid

precipitated from the mixture which had the same characteristics as the solid

formed in the manganese catalyzed autoxidation of benzaldehyde. Qualitative

examination of the reaction mixture after 90 per cent of the peracid decomposed

indicated that the amount of solid formed is proportional to the metal ion concen-

tration. No solid was observed in the uncatalyzed decomposition.

Paper chromatographic analysis of the soluble decomposition products

indicated that the only product formed was benzoic acid. The phenolic product

observed in the autoxidation reactions was not found in this case.

A. Order in perbenzoic acid

The logarithm of the measured perbenzoic acid concentration, log

BzOOH plotted against time gave a straight line as shown in Figures

18 and 19. This shows the decomposition to be first order in perbenzoic acid

for both the manganese (II) and (III) system.

B. Order in manganese ion

The rate of perbenzoic acid decomposition was followed for a series

of manganese (II) and (II) ion concentrations to establish the order in metal

ion. A plot of d log [BzOOH ] /dt versus catalyst concentration was linear as

shown in Figure 20. The decomposition of the peracid is thus first order in

manganese ions.

The decomposition rate may be expressed by:

-d [BzOOH] /dt = k [BzOOH] [Mn] (36)

where [Mn ] is the concentration of manganese (I) or (III). The value of the

rate constant, kj, for the manganese catalyzed decomposition of perbenzoic

acid at 500 is 5.9 1./mole/sec, for the manganese (II) system and 4.6 1./mole/sec.

for the manganese (Hi) system.

A small quantity of gas was evolved during the decomposition. Gas phase

chromatography indicated that the gas evolved was carbon dioxide. The amount

of carbon dioxide was found to be independent of the presence of oxygen over the

system. About one per cent of the perbenzoic acid decomposed to give carbon


It should be pointed out that the iodometric method used in following the

concentration of the peracid would not distinguish perbenzoic acid from any ben-

zoyl peroxide formed during the reaction.

O. OM Mn(II)

2.28 x 10-6 M

-3.2 -

g -3.4

-3.6 -

2.28 x 10-5 M

3.43 x 10-5 M

20 40 60 80 100 12

time (min. )

Figure 18. Decomposition of perbenzoic acid by manganese (II) acetate
( 1.09 x 10-2 Mperbenzoic acid, 500 )

.14 x 10-5 M



_ 0. 0 M Mn(III)

2.28 x 10-6 M






2.28 x 10-5 M

3.43 x 105 M


20 40 60

80 100 120

time ( min. )

Figure 19. Decomposition of perbenzoic acid by manganese (III) acetate
( 1. 09 x 10-2 M perbenzoic acid, 500 )

1.14 x 10-5 M








10 20 30 40

[Mn] x 106

Figure 20. Comparison of manganese (II) and manganese (III)
catalyzed decomposition of perbenzoic acid

m. Cobalt Systems

The results in this section show the effect of cobalt catalyst and

benzaldehyde concentration on the autoxidation at 500. In the concentration

range used the color of the cobalt catalyst was indiscernible. However, after

the oxidation was in process the reaction mixture turned a faint tan color.

The uptake of oxygen was nearly quantitative if the reaction was allowed

to approach completion. In no case was the difference between the amount of

oxygen used and the amount of benzaldehyde consumed over 2. 0 per cent.

The only products detected were benzoic acid and the intermediate

perbenzoic acid. Starting with 0.245 moles of benzaldehyde, 0.240 moles of

benzoic acid were recovered.

A. The effect of cobalt concentration

The dependence of the oxidation rate on the cobalt (1I) concentration was

examined by carrying out a series of runs at constant benzaldehyde concentra-

tion but varying the cobalt (II) ion concentration from 5.0 x 10-5 to 1.2 x 10-3

MyL(Table 12).

A plot of the oxidation rate against the square root of the cobalt concen-

tration, [ Co 1/2, was linear indicating a one-half order dependency on the

cobalt (II) concentration. This is shown in Figure 21.



( [BzH] o =0.988, 500, 1 atm. O2)

[Co(n)] d02/dt

M x 104 mM 02/min.

0.50 0.307

1.2 0.320

3.6 0.452

6.0 0.507

8.4 0.625

12.0 0.709




< 0.4



I I I i
1 2 3 4

[Co()] x 10o2

Figure 21. The effect of the cobalt concentration on the benzaldehyde
oxidation rate (0.988 M BzH, 500, 1 atm. 02 )


No induction periods or retardation effects were observed in these studies.

B. The effect of benzaldehyde concentration

A series of experiments at constant catalyst concentration (3.0 x 10-5

M cobalt) in which the bonzaldehyde concentration was varied between 0.1 and

1.0 M was performed to determine the order in aldehyde (Table 13).

A plot of the oxidation rate against the three-halves power of the benzalde-

hyde concentration was found to be linear as shown in Figure 22.

The rate law for tte cobalt catalyzed benzaldehyde autoxidation becomes:

rate = k [Co ] 1/2 [BzH ] 3/2 (37)

where [Co ] is the concentration of cobalt (II) ion and [ BzH ] is the concentra-

tion of benzaldehyde.

This is exactly the relationship found by Bawn (8) at 250.



([o(n)]= 3.0 x 10-5, 50, 1 atm. 02)

[BH] o -dO2/dt

mM 02/min.

0.197 0.98

0.396 0.272

0.593 0.360

0.988 0.581



- 0.4


I i I I I
2 4 6 8 10

[B.zH]/2 x 101

Figure 22. Variation of the oxidation rate with benzaldehyde concentration
( 3.0 x 10-5 MCo(III), 500, 1 atm. 2 )

IV. Cerium Systems

The effect of cerium salts on the autoxidation of benzaldehyde in acetic

acid was studied at 50. Cerium (IV) naphthenate and cerium (II) nitrate were

dissolved in acetic acid and these solutions were used in preparing the reaction


The solution was faintly yellow and turned to a straw color after oxida-

tion started. The only products isolated were perbenzoic acid, benzoic acid

and unreacted benzaldehyde. Figure 23 shows the course of a typical run. All

the benzaldehyde did not react. The extent of the reaction was found to be de-

pendent on the cerium concentration. This is shown in Table 14. Addition of

cobalt naphthenate (1.5 x 10-5 M) to the reaction after it had stopped caused a

rapid uptake of oxygen resulting in the complete oxidation of the benzaldehyde.

A. Induction periods

The effect of cerium ion concentration. A series of experiments in which

the concentration of Cerium (II) and (IV) was changed with the other parameters

held constant, showed the effect of the cerium concentration on the length of the

induction period. The results are shown in Table 14. The length of the induction

period was proportional to the cerium ion concentration up to 6 x 104 M, but the

length of the induction periods dropped off at higher concentrations, as shown in

Figure 23.



20 -

60 120 180 240 300 360 420 480

time (min. )

Figure 23. Typical cerium catalyzed run showing both inibition and
retardation (0.988 M BzH, 8.0 x 10- M Ce(IV), 500, 1 atm. 02 )



( [BzH = 0.988, 500, 1 atm. 02)

[Ce ] Oxidation d0/dt' ti Extent of reaction
M x 104 mM-O/min. mn. mM 02
-- 2 w.^--





















0. 279
























_ ~~ ____


The effect of the oxidation state of cerium. The effect of both the con-
centration of cerium (II) and cerium (IV) on the induction period is shown in

Figure 24. Both oxidation states gave nearly equal induction periods at equi-

valent concentrations.

B. Rate of oxidation

The effect of the cerium (iI) and (IV) concentrations on the post induc-

tion period oxidation rate of benzaldehyde is shown in Table 14. and plotted in

Figure 25. Low concentration of cerium ions gave a very slight increase in the

oxidation rate, but by and large, the concentrations of cerium (III) and (IV) ions

had little effect on the post induction period oxidation rate of benzaldehyde.

C. Extent of reaction

The effect of the cerium ion concentration on the extent of the reaction

is shown in Table 14. The extent of reaction was decreased as the concentra-

tion of cerium ion was increased for both oxidation states of the metal ion.

D. The effect of added initiator

The effect of a radical initiator on the cerium ion catalyzed system was

examined by adding 4.3 x 10-3 M 2,2, '-azobisisobutyronitrile to a reaction mix-

ture containing 0.988 M benzaldehyde and 4.0 x 10-4 M cerium (HI) and following

the oxidation at 50. The induction period was found to be 48 minutes for this

run compared to 73 minutes for an equivalent run without the added initiator.

The post induction period oxidation rate for the runs with and without initiator

were nearly the same.

0 Ce(IV)

o Ce(III)


2 4 6 8 10
[Ce] x 104

Figure 24. The effect of cerium ions on the induction period
(0.988 M BzH, 500, 1 atm. 02)



0 Ce(IV)

0.4 0 Ce(III)

0.3 0 O

ar- 0

Pe] x104

Figure 25. The effect of cerium ions on the post induction oxidation
rate (0.988 M BzH, 500, 1 atm. 02 )


The experimental results of this research are discussed on the

basis of current free radical theory. Particular attention is devoted to

the consideration of the role of manganese ions in benzaldehyde autoxi-

dation and to the manganese ion catalyzed decomposition of perbenzoic

acid. The effect of cobalt ions and cerium ions on the autoxidation are

considered and the different catalyst systems are compared.

I. Manganese System

The results of this research have shown that the manganese catalyzed

autoxidation of benzaldehyde is an extremely complex phenomena. Fortunately

the number of chemical species involved in this process is relatively small so

that a reasonable mechanism can be developed from the information obtained

in this work.

Prior to assembling the various bits of information into a mechanistic

picture, perhaps it would be advantageous to briefly summarize a few of the

important observed relationships. This summary can then be used as a guide

in developing the mechanism.


1. The reaction is first-order in benzaldehyde in the concentration

range 2 to 5 x 10-1 M and approaches zero-order as the benzaldehyde con-

centration is increased to 9 x 10-1 M.

2. The reaction is one-half order in either manganese (II) or (111) ions

at concentrations in the order of 1 x 10-5 M. However, as the metal ion con-

centration is further increased, the rate of oxygen uptake reaches a maximum

and then at still higher concentrations (1 x 10-3 M) decreases sharply. Induc-

tion periods appear at metal ion concentrations about 2 x 10-3 M.

3. The rate of oxidation is greater with manganese (III) than with

manganese (II).


4. The reaction is one-half order in perbenzoic acid over the concen-

tration range 1 to 5 x 10 M. At high initial concentrations, perbenzoic acid

inhibits oxygen uptake.

5. In the absence of benzaldehyde, perbenzoic acid is decomposed more

rapidly by manganese (II) than by manganese (III) ions. The rate law for the de-

composition is:

rate = k* [Mn ] [BzOOH] (36)

6. During an autoxidation the rate of benzaldehyde disappearance

parallels the rate of oxygen uptake and the rate of appearance of perbenzoic

acid. However, there is an abrupt halt of oxygen uptake and benzaldehyde con-

sumption which occurs when the perbenzoic acid concentration reaches a certain

point. After this the peracid concentration falls off sharply with no further oxy-

gen uptake.

It is apparent from these results that the path of the reaction is altered

easily and frequently.

A. Kinetics of manganese catalyzed benzaldehyde autoxidation

Since manganese ions play the central role in this process it might be well

to spell out this role. It has been pointed out by others (14, 39, 43, 45) and veri-

fied here that manganese (I) and (III) ions can decompose perbenzoic acid to

produce radicals. These radicals are capable of propagating the chain reaction.

Low catalyst concentration. At concentrations where a half-order depend-

ence of the rate on manganese ion concentration is observed, the chain sequence

is apparently:


Mnlu + BzOOH -- BzOO* or BzO* (38)

BzOO- + BzH Bz- + BzOOH (39)

Bz* + 02 --BzOO (40)

2 BzOO* ----inert products (41)

The rate law derived from this reaction sequence using the steady-

state approximation is:

dO2 kb 'p12
dt 71/2 BZH] [BzOOH] 1/2 [ Mn] 1/2 (42)

Results of this work are in complete agreement with this equation.

Thus a one-half order dependence on manganese ion, a first-order depend-

ence on benzaldehyde (at low benzaldehyde concentrations), a one-half order

dependence on perbenzoic acid and a zero-order oxygen dependence have been

observed. It thus appears that the reaction sequence given above adequately

explains the behavior of the system in this concentration range.

The value for the combined constants from equation 42., kp kbl/2 kt1/2,

for the manganese catalyzed autoxidation of benzaldehyde is 4.9.1/mole/sec.

High catalyst concentration. At higher manganese ion concentrations

the oxidation rate is decreased and induction periods are observed. From

this qualitative evidence it appears that manganese ions are able to play a

part in a radical termination process. This is not a unique idea; it has been

proposed as the scheme by which metal ions inhibit various radical chain reac-

tions including both autoxidations and polymerizations (54, 55, 56, 57,.68).


A reasonable chain sequence involving the metal ion in the termination

reaction is as follows:
Mn+n + BzOOH ---BzOO- or BzO. (43)

BzOO. + BzH ---Bz. + BzOOH (44)

Bz. + 02. BzOO. (45)

Mn+n + BzOO* ---Pinert products (46)

Derivation of the rate expression for the above reaction sequence is

accomplished with the aid of the steady-state assumption* and is:

1, [~EH] -kb [ Mn] (47)
kp kb
-dO2/dt=-- [BzH] [BzOOH] e

This equation may now be examined to see how well it fits the experi-

mental observations. The form of the equation is exactly that of the peak and

post peak portions of the curve in Figure 5. where the initial rate is plotted

against the manganese ion concentration at constant benzaldehyde concentra-

tion. It is in the range of manganese ion concentration represented by the

"falling-off" portion of the curve that inhibition by manganese ions is apparently

occurring. Recognizing that dO2/dt expresses the initial oxidation rate, so

that the benzaldehyde concentration is virtually constant, Equation 47. predicts

that the oxidation rate will be nearly independent of the manganese ion concen-

tration at low concentrations. In this range a change in the manganese concen-

tration will have little effect on the value of the exponential term. This

ee Appendix
See Appendix I.


corresponds to the "leveling-off" portion of the oxidation rate versus curve

(Figure 5). However as the manganese ion concentration is increased to

higher values it will have a significant effect on the exponential term and

- dO2/dt will be sharply decreased as is observed.

This is good evidence for the termination by manganese ions. Since

manganese (nI) and (II) ions exhibit the ability to decrease the rate of oxida-

tion it is likely that both are capable of terminating chains unless a mixture

of these ions is formed during the course of the reaction as postulated by the

Haber-Weiss mechanism.

Specific rate constants. Since Equation 48 apparently describes the

phenomenon occurring during the reaction, some information concerning the

rate constants kp, 1t, and kb can be obtained from the data. The values for

kb, obtained from the slope of a plot of In rate versus [ Mn ] for a series of

experiments summarized in Tables 2 and 3 are:

Ib Mn(II) =1.0 x 103 1./mole/sec.

kb Mn(II) = 5.1 x 102 1./mole/sec.

Similarly values for [BzH ] kb kp/ktt can be obtained from the intercepts of

these plots*. From these values the ratios of kp to kty are found to be:

Mn)= 8.0 x 10-3 tMn() = 1.85 x 10-2

Assuming Ic has the value reported by Ingles and Melville (7) ktt becomes:

A value for the initial perbenzoic acid concentration must be assumed
here. The value chosen is 1 x 10-5 M. This was selected because 1 x 10-4 M
was shown to increase the rate and iodometric analysis indicates that the con-
centration was not over 5 x 10-5 M.


ktMn(II) =4.1 x 105 1./mole/sec.

kt'Mn(I) = 1.7 x 105 1./mole/sec.

Thus the manganese ion termination rate constants have been estimated.

Comparison of rate constants. An interesting comparison can be made

between the kb values obtained in the manner described above and in those ob-

tained by following the rate of decomposition of perbenzoic acid in the presence

of manganese ion but in the absence of benzaldehyde (see Results, Section II).

The values obtained in the latter case are:

kbMn(II) =5.9 1./mole/sec.

kMn() = 4.6 1./mole/see.

These values are considerably lower than those reported above. It will be re-

called that the decomposition process which occurs in the absence of benzalde-

hyde is a non-chain reaction involving only initiation and termination of the

radicals, there being no species present for propagation. Furthermore the

analytical method used to follow the decomposition rate will not differentiate

between the peracid and benzoyl peroxide which is a highly probable product

conceivably formed by the reaction:

2 BzO. -- (BzO)2

As a result of this the kb values are useful only in considering relative rates.

Thus, a comparison of the ratio of I~ Mn(I) to kb Mn (mI)

with that of kb Mn(II) to kb Mn(II) reveals that the relative decomposition rates

are nearly identical. This suggests that the radical products from the manganese

(II) ion decomposition are the same as those from the manganese (II) decomposi-