**2. Classical mechanism of organic compounds liquid phase oxidation inhibited by stable nitroxide radicals**

For the first time Scheme 1 was proposed by A. Buchachenko with staff to describe the oxidation of ethylbenzene in presence of >NO• (II) and (III) (Buchachenko, 1965). *k*4/*k*<sup>1</sup> Values (Table 1) were obtained from kinetic data of some organic compounds' initiated oxidation in presence of >NO•. It can be seen that nitroxide radicals inhibit oxidation of methacrylic and acrylic ethers more effectively rather than oxidation of alkylaromatic compounds: *k*4/*k*1 values on average an order of magnitude greater for methacrylates and acrylates as compared with styrene, benzene, and cumene (Table 1).

At the same time *k*4/*k*1 value for methyl methacrylate (0.36) is close to that for cyclohexyl methyl ether (0.33) (Kovtun et al., 1974). Such results prove an important role of polar effects in reactions of >NO• with alkyl radicals during vinyl monomers oxidation.

So it should be noted that inhibiting activity of >NO• increases along with length of methacrylic ethers' alkyl substitute: *k*4/*k*1 value increases more than 3 times from methyl ether to ionylic ether. The similar trend also takes place for acrylic ethers (Table 1).


**Table 1.** *k*4/*k*1 Values in oxidizing monomers and hydrocarbons at 323 K

264 Nitroxides – Theory, Experiment and Applications

N O

N OH

>NO• (VII)

N O

OH

pyrroline, and imidazoline >NO• with active particles of chemical and biochemical

Structures of nitroxide radicals and corresponding hydroxylamines which reactions'

O O

N O

>NO• (I) >NOH (I) >NO• (II) >NOH (II) >NO• (III)

OC(O)C6H5

CN

CH2NH2

N O

>NO• (VIII)

>NO• (XI)

**2. Classical mechanism of organic compounds liquid phase oxidation** 

For the first time Scheme 1 was proposed by A. Buchachenko with staff to describe the oxidation of ethylbenzene in presence of >NO• (II) and (III) (Buchachenko, 1965). *k*4/*k*<sup>1</sup> Values (Table 1) were obtained from kinetic data of some organic compounds' initiated

N O

N O N OH

N OH

OC(O)C6H5

>NO• (V) >NO• (VI)

N

>NOH (V)

>NO• (IX)

N O N O

> N O

>NO• (X)

C6H11

N O

C6H5 O CONH2

>NO• (XII)

N O O

C6H5

OH

oxidation processes. Such an attempt has been made in this review.

mechanisms were analyzed in this work are presented in Figure 1.

N OH

N O

>NOH (VII)

CONH2

N OH CONH2

**Figure 1.** Structures of nitroxide radicals and corresponding hydroxylamines

>NOH (III) >NO• (IV)

>NOH (X)

**inhibited by stable nitroxide radicals** 

N OH

CONH2

NH2

Absolute *k*4 values were measured in (Aleksandrov et al., 1979) by ESR spectroscopy method (Table 2). We note that >NO• is one of the strongest acceptors for alkyl radicals. *k*4 Values are close to ones for molecular oxygen addition to alkyl radicals (*k*<sup>1</sup> ≥ 1∙107 M–1∙s–1 (Aleksandrov et al., 1979)). *k*4 Values for reactions of >NO• (I) – (VII) with alkyl radicals of methyl methacrylate at 323 K are within a limits of (0.8 – 2.0)∙107 M–1∙s–1 (Table 2). Even higher *k*<sup>4</sup> values for the reaction of >NO• (I) with R• of low molecular weight at 291 K were obtained in (Bowry & Ingold, 1992) by laser flash photolysis method: these values are within a limits of 1∙106 – 2∙109 M–1∙s–1.

The probability of >NO• participation in chain initiation process via its reaction with monomer's -bond was estimated in (Ruban et al., 1967) using the reaction of >NO• (I) and (III) with styrene and -methyl styrene as example. Reaction

>NO• + CH2=C(X)C6H5 >NO–CH2–C•(X)C6H5 (where X = H or CH3)

at 393 K proceeds with constant rate value which equals to 4.6 M–1∙s–1 in styrene, but in -methyl styrene this reaction doesn't proceed at all even at 453 K (Ruban et al., 1967). It's clear that >NO• initiating function completely suppresses by its participation in reaction (4).

Further we will see that reaction (4) is not the only one in oxidation inhibition by nitroxide radicals.


**Table 2.** Rate constants for the reaction M• (R•) + >NO• at 323 K (Aleksandrov et al., 1979)

## **3. Multiple chain-breaking by stable nitroxide radicals**

It was proved on oxidation of a number of compounds that oxidation chains propagate by peroxide radicals which possess redox properties. These are HO2• radicals (cyclohexadiene (Howard & Ingold, 1967), 1,2-ethylene substituted and 1,4-butadiene substituted monomers (Mogilevich & Pliss, 1990)), >C(OH)O2• (alcohols (Kharitonov & Denisov, 1967)), and >CH– CH(OO•)N< (aliphatic amines (Aleksandrov, 1987)). Dual reactivity of these radicals results in multiple >NO• participation in chain termination processes (Denisov, 1996). So for hydroperoxide radical this process can be described with the following reactions (Denisov, 1996):

$$\mathrm{HO}\bullet\mathrm{^{\bullet}} \mathrm{+} \mathrm{NO}\bullet\longrightarrow\mathrm{^{\bullet}} \mathrm{NOH} \mathrm{+} \mathrm{O}\bullet\tag{5.1}$$

$$\text{HO} \bullet \text{'} + \text{NOH} \longrightarrow \text{NO} \bullet \text{'} + \text{HO} \text{'} \tag{5.2}$$

Let's perform the analysis of oxidation mechanism with >NO• regeneration and one without it.

### **3.1. Analysis of oxidation mechanism without nitroxide radical regeneration**

In accordance with Scheme 1 initial rate process (*W*) without >NO• regeneration would be described by the following equation (*W* = *W*0 when [>NO•]0 = 0):

$$\mathbf{W}\_{i}\left(\frac{\mathbf{W}\_{0}}{\mathbf{W}} - \frac{\mathbf{W}}{\mathbf{W}\_{0}}\right) = \frac{k\_{4}\text{[}\text{>NO}^{\bullet}\text{]}\_{0}\mathbf{W}\_{0}}{k\_{1}\text{[}\text{O}\_{2}\text{]}}\tag{1}$$

If *k*4[>NO•]0 >> *k*3[RO2•] then

266 Nitroxides – Theory, Experiment and Applications

radicals.

1996):

Further we will see that reaction (4) is not the only one in oxidation inhibition by nitroxide

~CH2C•HC6H5 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COOCH3 ~CH2C•(CH3)COO–*n*-C4H9 ~CH2C•(CH3)COO–*iso*-C4H9 ~CH2C•HCOOCH3 ~CH2C•HCOO–*iso*-C4H9 *cyclo*-C•6H11 (CH3)2C•CN (in benzene)

**Table 2.** Rate constants for the reaction M• (R•) + >NO• at 323 K (Aleksandrov et al., 1979)

It was proved on oxidation of a number of compounds that oxidation chains propagate by peroxide radicals which possess redox properties. These are HO2• radicals (cyclohexadiene (Howard & Ingold, 1967), 1,2-ethylene substituted and 1,4-butadiene substituted monomers (Mogilevich & Pliss, 1990)), >C(OH)O2• (alcohols (Kharitonov & Denisov, 1967)), and >CH– CH(OO•)N< (aliphatic amines (Aleksandrov, 1987)). Dual reactivity of these radicals results in multiple >NO• participation in chain termination processes (Denisov, 1996). So for hydroperoxide radical this process can be described with the following reactions (Denisov,

HO2• + >NO• >NOH + O2 (5.1)

HO2• + >NOH >NO• + H2O2 (5.2)

Let's perform the analysis of oxidation mechanism with >NO• regeneration and one without it.

In accordance with Scheme 1 initial rate process (*W*) without >NO• regeneration would be

[O ] *<sup>i</sup> W kW <sup>W</sup> <sup>W</sup> WW k*

 

0 4 00 0 12

•

(1)

[>NO ]

**3.1. Analysis of oxidation mechanism without nitroxide radical regeneration** 

described by the following equation (*W* = *W*0 when [>NO•]0 = 0):

**3. Multiple chain-breaking by stable nitroxide radicals** 

M• >NO• *k*4, M–1∙s–1

I I II III V VI VII VIII I I I I I I

8.0∙106 1.2∙107 0.8∙107 0.8∙107 1.4∙107 1.6∙107 1.2∙107 2.0∙107 1.2∙107 0.9∙107 3.0∙107 1.8∙107 3.0∙107 8.6∙107

$$\mathcal{W} = \frac{k\_1[\mathbf{O}\_2]}{k\_4[\text{!}\text{!}\text{NO}^\bullet\text{]}\_0} \mathcal{W}\_i \tag{2}$$

So the rate process in linear termination mode (high [>NO•]0) is directly proportional to the partial oxygen pressure (*P*o2). Therefore if oxygen is substituted by air the oxidation rate is to decrease 5 times. Such results were gained in (Browlie & Ingold, 1967; Pliss & Aleksandrov, 1977). But here is one important circumstance. From literature data (Aleksandrov, 1987) it's known that the reduction of nitroxide to corresponding hydroxylamine (>NOH) occurs via the reaction of aminoalkyl radical >N–C•H–CH< with nitroxide radical as >NO• attack to -C–H bond of alkyl radical.

$$\text{\textbullet N-C-H-CH-} + \text{NO}^\* \longrightarrow \text{\textbullet NOH} + \text{\textbullet N-CH=} < 1$$

The hydroxylamines being formed are thermally stable under experimental conditions (Aleksandrov, 1987). If we assume that >NOH is able to react with R•

>N–C•H–CH< + >NOH >N–CH2–CH< + >NO•,

then >NO• stoichiometric coefficient must be more than 1.

Let's consider the probability of >NO• cross-dispropotionation with other alkyl radicals. Such a consideration is quite useful cause at physiological *P*o2 values in body tissues of higher animals and humans (5 – 50 torus) oxygen concentration is less than 1∙10–4 M (Porter & Wujek, 1984). In this case inequality [RO2•] >> [R•] is not satisfied and it's necessary to take into account alkyl radicals' participation contribution. This can be done by modifying Scheme 1 for vinyl monomers' oxidation (Scheme 2).

> (i) Initiator O ,M <sup>2</sup> M• *Wi*

(1) M• + O2 MO2• *k*<sup>1</sup>

(2) MO2• + M M• *k*<sup>2</sup>

(3) MO2• + MO2• molecular products *k*<sup>3</sup>


**Scheme 2.** Mechanism of organic compounds oxidation inhibited by nitroxide radicals taking into account M• with >NO• disproportionation

Let's estimate >NO• recombination and disproportionation shares ratio according to reactions (4.1) and (4.2). There's almost no any experimental data for such estimation, so we have to use the results of quantum-chemical calculations.

Table 3 represents the values of quantum energies of >NO• reactions1 with peroxyalkyl radicals ~OOM• (DFT B3LYP/6-31G\* calculation similar to one in (Becke, 1993)). As a structural unit we've used –OOCH3 fragment. As can be seen from the table, such operation is quite acceptable: substitution of –OOCH3 to –CH2CH3 or to –CH2CH2CH3 doesn't result in significant changes in energy values calculated.

It should be noted that in accordance with calculated results recombination's probability is significantly greater as compared with disproportionation's one: mean difference in energies is greater or equal of 30 kJ/mol. However, cross-disproportionation in liquid phase also can not be excluded: polar effects may have a significant effect especially if there are polar groups in conjugation with -C–H bond (Roginskii, 1987). Recombination and disproportionation energies lowering during methyl group addition to -position of radical center ~OOM• also seem to be a logical cause as a steric effects appear in this case.


**Table 3.** Reaction energies of alkyl radicals with >NO• (I) (kJ/mol)

Reactions (4.2) and (4.3) rates ratio can be estimated on the basis of experimental kinetic data of >NO• (I) consumption in cumene, styrene, or methyl methacrylate in inert atmosphere. Experiment conditions: 323 K, atmosphere of argon, [>NO• (I)]0 = 5∙10–3 M, initiator – azobisisobutyronitrile, *Wi* = 1∙10–7 M∙s–1. The dynamic equilibrium is to set over time in case of reactions (4.2) and (4.3) and observed residual ESR signal value is to increase along with [>NO•]0 growth since the following equalities are valid during that equilibrium:

<sup>1</sup> In chemical thermodynamics quantum-chemically calculated reaction energy is used. It is a difference between full energies of reaction's products and reagents. This value often correlates with experimental value – enthalpy of reaction (http://cccbdb.nist.gov/).

Kinetics and Mechanism of Reactions

$$\begin{aligned} k\_{4,2}[\mathbf{R}^\bullet][\overline{\rhd \mathbf{NO}^\bullet}] &= k\_{4,3}[\mathbf{R}^\bullet][\overline{\rhd \mathbf{NO}^\bullet}] \\\\ [\overline{\rhd \mathbf{NO}^\bullet}] &= [\rhd \mathbf{NO}^\bullet]\_0 - [\overline{\rhd \mathbf{NO}^\bullet}] \\\\ [\overline{\rhd \mathbf{NO}^\bullet}] &= \frac{k\_{4,3}[\mathbf{>} \mathbf{NO}^\bullet]\_0}{k\_{4,2} + k\_{4,3}}, \text{ whence } \frac{k\_{4,2}}{k\_{4,3}} = \frac{[\rhd \mathbf{NO}^\bullet]\_0}{[\overline{\rhd \mathbf{NO}^\bullet}]} - 1 \end{aligned}$$

After completely >NO• consumption the residual ESR signal amplitude doesn't exceed a noise level under experimental conditions. This corresponds to potential stationary >NO• concentration of less than 10–7 M (ESR spectrometer Adani CMS 8400). In this case [>NOH] [>NO•]0 and at [>NO•]0 = 5∙10–4 M we have the following value of *k*4.2/*k*4.3 ratio:

*k*4.2/*k*4.3 = (5∙10–4/1∙10–7) – 1 5∙104

And now the ratio of (4.3) to (4.2) reaction rates can be estimated:

268 Nitroxides – Theory, Experiment and Applications

significant changes in energy values calculated.

Table 3 represents the values of quantum energies of >NO• reactions1 with peroxyalkyl radicals ~OOM• (DFT B3LYP/6-31G\* calculation similar to one in (Becke, 1993)). As a structural unit we've used –OOCH3 fragment. As can be seen from the table, such operation is quite acceptable: substitution of –OOCH3 to –CH2CH3 or to –CH2CH2CH3 doesn't result in

It should be noted that in accordance with calculated results recombination's probability is significantly greater as compared with disproportionation's one: mean difference in energies is greater or equal of 30 kJ/mol. However, cross-disproportionation in liquid phase also can not be excluded: polar effects may have a significant effect especially if there are polar groups in conjugation with -C–H bond (Roginskii, 1987). Recombination and disproportionation energies lowering during methyl group addition to -position of radical

center ~OOM• also seem to be a logical cause as a steric effects appear in this case.

<sup>M</sup>• Reaction type

C6H5CH•CH2OOCH3 –121 –77 C6H5CH•CH2CH3 –136 –69 C6H5CH•CH2CH2CH3 –125 –70 C6H5C•(CH3)CH2OOCH3 –79 –70 C6H5C•(CH3)CH2CH3 –94 –68 C6H5C•(CH3)CH2CH2CH3 –69 –67 CH3OC(=O)CH•CH2OOCH3 –127 –111 CH3OC(=O)CH•CH2CH3 –172 –145 CH3OC(=O)CH•CH2CH2CH3 –144 –93 CH3OC(=O)C•(CH3)CH2OOCH3 –79 –61 CH3OC(=O)C•(CH3)CH2CH3 –93 –75 CH3OC(=O)C•(CH3)CH2CH2CH3 –88 –79

**Mean value –111 –82** 

Reactions (4.2) and (4.3) rates ratio can be estimated on the basis of experimental kinetic data of >NO• (I) consumption in cumene, styrene, or methyl methacrylate in inert atmosphere. Experiment conditions: 323 K, atmosphere of argon, [>NO• (I)]0 = 5∙10–3 M, initiator – azobisisobutyronitrile, *Wi* = 1∙10–7 M∙s–1. The dynamic equilibrium is to set over time in case of reactions (4.2) and (4.3) and observed residual ESR signal value is to increase along with

<sup>1</sup> In chemical thermodynamics quantum-chemically calculated reaction energy is used. It is a difference between full energies of reaction's products and reagents. This value often correlates with experimental value – enthalpy of reaction

[>NO•]0 growth since the following equalities are valid during that equilibrium:

**Table 3.** Reaction energies of alkyl radicals with >NO• (I) (kJ/mol)

(http://cccbdb.nist.gov/).

M• + >NO• MON< M• + >NO• M–H + >NOH

$$\frac{W\_{4.3}}{W\_{4.2}} = \frac{k\_{4.3} \overline{[> \text{NOH}]}}{k\_{4.2} \overline{[> \text{NO}^{\bullet}]}} \approx 2 \cdot 10^{-5} \frac{\overline{[> \text{NOH}]}}{[> \text{NO}^{\bullet}]} \cdot 1$$

Therefore in range of up to 99% of >NO• consumption reaction's (4.3) share is less than 1% of reaction (4.2), so practically there's no any >NO• regeneration at all. That is *f* = 1 in inert atmosphere and in these substrates' medium. It's obvious that reaction (4.2) would be completely suppressed by reaction of >NOH with RO2• at [O2] > 1∙10–4 M.

#### **3.2. Analysis of oxidation mechanism with nitroxide radical regeneration**

In case of >NO• regeneration the oxidation's mechanism describes by Scheme 3.


**Scheme 3.** Mechanism of organic compounds oxidation inhibited by nitroxide radicals taking into account >NO• regeneration

The following equation is valid for this scheme:

$$\mathcal{W}\_{\text{i}} \left( \frac{\mathcal{W}\_{0}}{\mathcal{W}} - \frac{\mathcal{W}}{\mathcal{W}\_{0}} \right) = \frac{k\_{\text{4}} \text{[} \text{>} \text{NO}^{\bullet}\text{]}\_{0} \mathcal{W}\_{0}}{k\_{\text{1}} \text{[} \text{O}\_{2}\text{]}} + \frac{2k\_{\text{5}} \text{[} \text{>} \text{NO}^{\bullet}\text{]}\_{0} \mathcal{W}\_{\text{i}}^{0.5}}{k\_{\text{3}}^{0.5}} \tag{3}$$

where *k*5 = (*k*5.1[>NO•] + *k*5.2[>NOH])/2[>NO•]0.

Kinetic analysis shows that at [O2] ~ 1∙10–2 M and [>NO•]0 < 10–4 M the contribution of reaction (4) to chain termination process is negligible, and then

$$\frac{\partial W\_0}{\partial W} - \frac{W}{W\_0} = \frac{2k\_5 \text{[>NO}^\bullet\text{]}\_0}{\left(W\_i k\_3\right)^{0.5}}\tag{4}$$

With the drop of *P*o2 and small share of quadratic chain termination the oxidation rate will decrease not linearly, but slower. Such facts were found for instance in (Pliss & Aleksandrov, 1977; Ruban et al., 1967).

Reaction (5.1) proceeds as disproportionation of nitroxide and peroxide radicals (Denisov, 1996):

$$\begin{aligned} \text{HO} \bullet ^\bullet + \text{>NO}^\bullet &\longrightarrow \text{>NOH} + \text{Oz} \\ \rhd \text{C(OH)} \bullet ^\bullet + \text{>NO}^\bullet &\longrightarrow \text{>NOH} + \text{>C=O} + \text{Oz} \\ \rhd \text{CH-CH(OO^\bullet)} \text{N} \longleftrightarrow \text{>NO}^\bullet &\longrightarrow \text{>NOH} + \text{>C=CH-N} + \text{Oz} \end{aligned}$$

>NO• regeneration and multiple chain termination processes are caused just by subsequent reaction (5.2). Measured kinetic inhibiting factors2 for different nitroxide radicals and substrates presented for example in review (Denisov, 1996). The most of *f* values greater than ten and reflects just a lower bound of this value.
