Oxidation of Organic Compounds

9 Ionic Catalysis in Chain Oxidation of

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Alcohols E. T. DENISOV, V. M. SOLYANIKOV, and A. L. A L E X A N D R O V Institute of Chemical Physics, Academy of Sciences, Vorobyevskoye Chaussee 2-b, Moscow, USSR H O decomposes to free radicals in 2-propanol by the action of H . Free radicals are also produced by the reaction between tert-BuOOH and Br in 1-propanol. The H C O ions inhibit the oxidation of cyclohexanol initiated by AIBN, destroying many oxyperoxide radicals—i.e., H C O is a negative catalyst. Appropriate reaction schemes and rate equations are proposed. 2

2

+

-

3

3

-

-

/âtalysis by transition metals in liquid-phase oxidation has been thor^ oughly investigated. The roles of other ions have not been sufficiently studied. This paper is concerned with catalysis by hydrogen ions and some anions, i n the chain oxidation of secondary alcohols such as cyclohexanol and 2-propanol. Secondary alcohols, because of their polarity, are convenient for studying ionic homolytic reactions and their role in chain oxidation. The oxidation of secondary alcohols by oxygen may be represented by the following chain reaction mechanism. Initiator

>C-

/OH

.

r

1 1

+ o2

,

>C-

>c:

,OH

>c:

.OH OO

,OH

ÔH > < OOH

+ >c

.OH

.OH

OO

"OOH

+

ÔH >C-

+ Oo + >c=o

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

9.

113

Oxidation of Alcohols

DENISOV E T A L .

The rate of the reaction is R =

(1)

k kf* p

For cyclohexanol, k k f — 5.0 Χ 10 exp ( - 10,800/RT) 1. molesec.- , k = 1.1 χ 10 exp ( - 11,900/ΚΓ) 1. mole" sec. , k — 5.0 X 10 exp X ( - 2200/RT) ( I ) . For 2-propanol, k k f = 3.0 X 10 exp ( - 12,000/RT) 1. m o l e ' sec.- ( 2 ) . 1/2

3

1 / 2

p

1/2

7

p

1

6


1

2

t

4

1 / 2

p

1/2

1

1/2

1/2

Hemolytic Decomposition of Hydrogen Peroxide by Action of Hydrogen Ions Strong mineral acids such as H S 0 , H C 1 0 , H C I , and H N 0 were found to accelerate the chain oxidation of 2-propanol. Experiments on 2-propanol oxidation with an A I B N initiator showed that acids do not affect the reactions of peroxide radicals since they do not change the value of k k f . Oxidation is accelerated because of the reaction between H 0 and H involving the generation of free radicals, 2

p

2

4

4

3

1 / 2

+

2

2

4

[H 0 ] XI0 2

2

2

4

MOLES/LITER

Figure 1. Rehtion of R to square of concentra tion of H 0 in 2-propanol at 70°C. with 0.0925M H SO A

2

2

2

h

The 2-propanol was oxidized at 70°C. The rate of oxygen consump tion was measured manometrically. The rate of free radical generation, R was calculated from Expression 1 (at 70°C. k k f = 1.5 10" 1. mole" sec.~ ). In the absence of acid, free radicals were formed in the reaction between H 0 and 2-propanol (2), k = RJ[H 0 ] = 5 Χ 10" sec." (at 7 0 ° C ) ; in the presence of 1.8 Χ 10" mole per liter H C 1 0 , h = 6.3 Χ 10" sec." ( at 70°C. ). It was found that R* ~ [ H 0 ] ( Figure 1 ). i?

p

1/2

3

1 / 2

1/2

1/2

2

1

2

{

2

8

2

2

6

1

4

2

2

2

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

114

OXIDATION

OF

ORGANIC

COMPOUNDS

1

Ion concentration in the 2 - p r o p a n o l - H S 0 system was measured by the electroconductivity method and calculated from the expression [ion] = #c/A , where κ is the specific electroconductivity, A is the equivalent electroconductivity, and A = 61 sq. cm. ohm" m o l e (70°C. ). 2

4

x

x

1


x

-1

_J

I

I

I

I

2

4

6

8

10

+

[H ]

3

ΧΙΟ MOLES/LITER

Figure 2. Relation of R i to concentra tion of H in 2-propanol at 70°C. with 0.0181 mole/liter H 0 . Estimated by electroconductivity method +

2

2

The rate of initiation increases linearly with the concentration of hydrogen ions ( Figure 2 ). It decreases when a base ( pyridine ) is intro duced into the system 2-propanol + hydrogen peroxide + perchloric acid + oxygen. A l l the above is evidence that the reaction takes place between hydrogen peroxide and hydrogen ions. The rate of initiation is Ri = k [ H 0 ] [ H ] . For sulfuric acid, k = 0.14 sq. liters mole" sec." (70°C. ). The energy of activation, E, is 27 ± 2 kcal. per mole. The total decomposition of hydrogen peroxide by the action of hydrogen ions was measured iodometrically in the absence of oxygen. The rate of hydrogen peroxide decomposition was found to be 2

2

2

+

2

1

R = * [ H 0 ] [H ] 2

+

2

In the case of perchloric acid Ε = 24.4 ± 2 kcal. per mole; k = 1.03 X 10 exp ( - 2 4 400/RT) liters m o l e sec." Comparing R and Ri shows that decomposition of hydrogen peroxide into free radicals represents only a small part (from 1 to 5 % ) of the over-all decomposition. It follows that the greater part of hydrogen 14

1

1


9.

115


ET A L .

DENisov

peroxide decomposes heterolytically under the action of acids. mechanisms may be proposed to explain the experimental results : HA + ROH ^ A" + R O H Κ R O H + H 0 ^ ROH + H 0 2

2

( 3


2

2

3

+

2

+

*i k l

\ H 0 I /

2

+

H 0 + OH -> H 0 + O H 2

2

TT

O

\ OH + >CC^ +

OH + H 0 . H 0 +

2

^2

H

H

+ +

H 0 + H + >C=0 +

2

i

k +

=

3

fc [ROH] 2

- HO- + H + H 0 +

2

(

+

Two

2

:

3

H 0 3

2

2

+ H 0

+

2

2

+ ROH

+

-» H 0 + HO- + H 0 +

3

^5

2

H 0 + H 0 + >C=0 3

+

2

k±K fc + fc [ROH] 4

5

If O H is formed in the system, R O O H would be expected to form by the reaction: +

H O 4- R O H ^ R O O H + H +

+

It was found by iodometric analysis that hydroperoxide was not produced in the system R O H + H 0 + H . Consequently, Mechanism II seems to be more probable. 2

2

+

Similar results were obtained for tert-butyl hydroperoxide and per chloric acid in 2-propanol. Thus, it is evident from the decomposition of hydrogen peroxide into free radicals that both heterolytic and homolytic reactions may be catalyzed by hydrogen ions. Further research is needed to investigate proton catalysis in certain homolytic reactions. Initiation by Reaction between ROOH and Br" Initiation by hydroperoxides may be intensified by anions capable of reducing other compounds. The formation of radicals in the reaction between ferf-butyl hydroperoxide and tetraethylammonium bromide was studied at 55° to 77 ° C in the presence of oxygen with 1-propanol as solvent. The rate of initiation was measured by consumption of a-naphthol that acted as acceptor of free radicals. Naphthol was determined colorimetrically. Two radicals were suggested to react with one molecule of α-naphthol. The rate of initiation was found to be (Figure 3). Ri = fc [ROOH] + fc[ROOH] [ E t N B r ] 0

4



116

OXIDATION OF ORGANIC COMPOUNDS

I

0 Figure 3.

1

2

R+ dependence on concentrations at 70°C.

1. tert-BuOOH at [EuNBr] = 0.1 M 2. EttNBr at [ROOH] = 0.056M 3. Br at [ROOH] = 0.056M

The degree of E t N B r dissociation into ions i n 1-propanol was measured b y the electroconductivity method. The rate of initiation was found to depend linearly on both [ E t N B r ] and [ B r ] . Experiments with a 1-propanol-water-hydroperoxide system at constant concentration of E t N B r were carried out to determine whether E t N B r or Br" reacted with the hydroperoxide. Water increased the degree of E t N B r dissociation. The rate of initiation was found to increase linearly with Br" at a constant concentration of E t N B r . Thus, the reaction seems to take place between hydroperoxide and Br". The following mechanism may be proposed: 4

4

4

4

4

4

R O O H + Br" - » RO- + H O ' + BrThe rate constant is k = 4.0X 10 exp(-19500/RT) liters mole" sec." 8

1

1


9.

DENisov E T A L .

117


Negative Catalysis by H C 0 ~ Ions in Chain Oxidation of Cyclohexanol 3

The inhibiting effect of N a H C 0 on chain oxidation was established by studying the effect of ions on the oxidation of cyclohexanol. The latter was oxidized at 75°C. with A I B N as initiator (R, = 6.9 Χ 10 mole l i t e r sec." ). T o dissolve N a H C 0 , 9 % of water was added to cyclohexanol. The rate of oxidation was measured volumetrically. 3

7


1

-1

3

MINUTES Figure 4. Kinetic curves of oxygen con sumption in cyclohexanol oxidation at 75°C with 0.01M AIBN Ri = 6.9 X 10~ mole/liter/second 1. Without inhibitor 2. With 1 X 10~ M 1-naphthol 3. With 4 X WW NaHCO 7

4

s

The constancy of the low oxidation rate i n the presence of H C O ~ is surprising. As seen from Figure 4, the ordinary inhibitor of reaction, α-naphthol i n a concentration of 10" mole per liter, disappears i n a few minutes, and the reaction is rapidly accelerated. The H C 0 " ion ( i n 4 X 10" mole per liter concentration) inhibits the oxidation of cyclohexanol for half an hour, and the rate of reaction does not increase during this period. Thus, the H C 0 " ion inhibits the chain oxidation not as an ordinary inhibitor but as a negative homogeneous catalyst, and each ion of such a catalyst may terminate many chains. H

4

3

5

3

The peroxy radicals of secondary alcohols are oxyperoxide radicals. Experiments using various systems show the relationship between the structure of peroxide radicals and negative catalysis by H C 0 ~ ( Table 1 ). 3


118

OXIDATION OF ORGANIC COMPOUNDS

Table I.

1

Effect of N a H C 0 on Oxidation of Various Systems 3

75°C; 0.01 mole per liter A I B N ; [ N a H C 0 ] = 1.4 Χ 10" mole per liter R, Rate in presence of N a H C 0 R , Rate in absence of N a H C 0 I. Cyclohexanol II. Cyclohexanone 4

3

3

3

0

Structure of Peroxide Radicals


System

9

S!i

! 0

> C

89.3% II 5.7% H 0 5.0% I 2

-

Ν

2

2

H

74.3% II 5.7% H 0 20.0% tert-BuOH

>C'

οο·

\ ~ '

2

1

R/R

1

0

H

X

1

6

0

ÔO-

>C^

94.3% II 5.7% H 0 0.15 mole liter" H 0

R ,10~ Mole Liter' Sec.'

H

4

Η

OH >C

4

0

6

4

>C—O-

οο·

2.92

0.61

3.00

1.00

ôo

H

>C^ ν

7

^οο·

00·

2

0

H C 0 " ions inhibit chain oxidation only when oxyperoxide radicals appear in the system. The experimental dependence of oxidation rate on H C 0 " (Figure 5) at 75°C. is R 2/R2 ι + [HC0 ], a = 1.8 Χ 10 liters per mole 3

3

O

=

fl

5

3

The results obtained may be interpreted by using the following scheme: .OH >c

ôo-

+ H C

° "

ÔH

> -> c

00

k,

3

.OH

> < —O-O H

+

È

° 3

^

+co 3

2—

> C = O + O + HCO 2

3

+ HCO,"


9.

DENisov

ET AL.

Oxidation

of

119

Alcohols


•

Figure 5.

Dependence of rate of cyclohexanol oxidation with AIBN on concentration of HCOf Ri = 6.9 X 10' mole/liter/second 1. At 75° C. 2. At 65°C. 3. Plot of R o / R against HCOf • 65°C. Ο 75°C. 7

2

2

B y assuming an equality M H C 0 ] [R0 ] = fc [>CHOH] [ C 0 ] it is possible to obtain for small H C 0 " ( f c [ > C H O H ] > k [HC0 ''] ) 3

Q R

_-,

2

3

p

3

R

3

2

,

1

_

3

fci*2[HCQ -] [ > C H O H ] 8

^

2

T

By comparison with the experimental function, we obtained

(at 7 5 ° C , 9 % H 0 ) , [ > C H O H ] = 8.75 moles per liter, k = 2.1 χ 10 liters mole" sec." . Therefore, fcifc A3 = 3.3 Χ 10 liters mole" sec." . Thus, certain anions may be seen to serve as negative catalysts in the chain oxidation of alcohols. 2

1

δ

t

1

11

2

1

1

Literature Cited (1) Alexandrov, A. L., Denisov, E . T., Nauk

1966,

Izv. Akad.

1737.

(2) Solyanikov, V. M., Denisov, E . T., (1964).

Neftekhimia

Nauk

SSSR,

Otd.

Khim.

3, 360 (1963); 4, 458

RECEIVED November 16, 1967.


Oxidation of Organic Compounds

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