Rate of Oxidation of Some Amines and Ammonium Compounds by Potassium Permanganate

V. I. Lutkova

In previous works on determining the rate of oxidation of organic compounds it was noted that, for certain amines, the rate of interaction with permanganate depends especially strongly on whether the reaction medium is neutral, acidic, or alkaline. Thus, amines of the saturated series react very slowly with potassium permanganate in the presence of acids [4] and considerably faster in an alkaline medium; as the table below shows, the difference reaches \(10^6\) (for methylamine). For pyridine, on the contrary, the rate of oxidation in alkaline solution is very low, while in neutral or acidic solution it increases sharply [5]. In some cases the influence of the medium is insignificant—such are the amines of the aromatic series [6].

Table 1

As Table 1 shows, the rate of oxidation as a function of the reaction medium varies very considerably, depending on the nature of the amine, and, evidently, must be governed by several causes. Among the structural and electronic factors, the following may affect the rate of the reaction:

1. The electrical state of the substance being oxidized and of the oxidizing agent. In the oxidation–reduction process, electrons pass from one substance (the reducing agent) to another (the oxidizing agent). The ease of this transfer, of course, must depend on whether the reacting components are present as neutral molecules or as ions of one sign or the other. It is clear that one and the same group of atoms should give up electrons more readily when it has a negative charge than when it is in the state of a neutral molecule. Thus, the \(\mathrm{SO_3}\) group may be a molecule (sulfurous anhydride) or an anion of sulfurous acid and its salts:

\[ \mathrm{Na_2SO_3 \rightleftarrows 2Na^{\cdot} + SO_3^{\prime\prime}.} \]

In the form of a molecule this group is an oxidizing agent (although a rather weak one); the sulfite ion represents a fairly strong reducing agent.

The iron ion may be divalent or trivalent:

\[ \mathrm{FeCl_2 \rightleftarrows Fe^{\bullet\bullet} + 2Cl'} \]

\[ \mathrm{FeCl_3 \rightleftarrows Fe^{\bullet\bullet\bullet} + 3Cl'} \]

Divalent iron is a reducing agent, trivalent iron an oxidizing agent. In the oxidation of \(\mathrm{Fe^{\bullet\bullet}}\) to \(\mathrm{Fe^{\bullet\bullet\bullet}}\), an electron is lost, and it turns out that a neutral molecule loses an electron more readily than a positive ion. Therefore, strongly dissociating ferrous chloride is oxidized with more difficulty than the weakly dissociating hydroxide \(\mathrm{Fe(OH)_2}\). With trivalent iron the reverse is observed. The oxidizing action of \(\mathrm{Fe^{\bullet\bullet\bullet}}\) is considerably greater in strongly dissociating salts than in the hydrate of the oxide \(\mathrm{Fe(OH)_3}\). The cation \(\mathrm{Fe^{\bullet\bullet\bullet}}\) attracts electrons more strongly than a neutral molecule can. An amine may be a neutral molecule or may give a cation

\[ \mathrm{CH_3NH_2 \quad or \quad CH_3NH_3Cl \rightleftarrows (CH_3NH_3)^{\bullet} + Cl'}. \]

In the form of a cation it must, of course, lose electrons with greater difficulty, i.e., be oxidized. This can explain the sharp slowing of the reaction with permanganate in the presence of acids. However, pyridine in the form of salts is oxidized more readily.

1. Valence of nitrogen. The best studied is the transition of this element from the trivalent state to the pentavalent state and back. It is known that an increase in the valence of nitrogen weakens the bonds of its atom. This explains, for example, the acceleration of the hydrolysis of amides in an acidic medium:

\[ \begin{array}{ccc} \mathrm{CH_3} & \mathrm{CH_3} & \mathrm{CH_3} \\ | & | & | \\ \mathrm{C{=}O + HOH} & \rightleftarrows \mathrm{C}\!\begin{matrix}/\mathrm{OH}\\ \backslash\mathrm{OH}\end{matrix} & \rightarrow \mathrm{C{=}O + NH_3} \\ | & | & | \\ \mathrm{NH_2} & \mathrm{NH_2} & \mathrm{OH} \end{array} \]

the reaction proceeds slowly,

\[ \begin{array}{ccc} \mathrm{CH_3} & \mathrm{CH_3} & \mathrm{CH_3} \\ | & | & | \\ \mathrm{C{=}O + HOH + HCl} & \rightarrow \mathrm{C}\!\begin{matrix}/\mathrm{OH}\\ \backslash\mathrm{OH}\end{matrix} & \rightarrow \mathrm{C{=}O + NH_4Cl} \\ | & | & | \\ \mathrm{NH_2} & \mathrm{NH_3Cl} & \mathrm{OH} \end{array} \]

acceleration by many times.

This same weakening of bonds is used in the synthesis of secondary amines from primary ones, by heating the latter with their salts:

\[ \mathrm{C_6H_5NH_3Cl + HNHC_6H_5 \rightarrow C_6H_5NH{-}C_6H_5 + NH_4Cl.} \]

The benzidine rearrangement, which proceeds under the influence of acid, belongs here as well:

\[ \mathrm{C_6H_5NHNHC_6H_5 + 2HCl \rightarrow C_6H_5NH{-}NH{-}C_6H_5 \rightarrow H_2N{-}} \]

\[ \begin{array}{ccccc} & \diagdown\!\!\diagup & \diagdown\!\!\diagup & & \diagdown\!\!\diagup \\ & \mathrm{H}\quad \mathrm{Cl} & \mathrm{H}\quad \mathrm{Cl} & & \mathrm{H}\quad \mathrm{Cl} \end{array} \]

\[ \mathrm{{-}C_6H_4{-}C_6H_4NH_2,} \]

\[ \begin{array}{c} \diagdown\!\!\diagup \\ \mathrm{H}\quad \mathrm{Cl} \end{array} \]

It is possible that this same circumstance contributes to the easier oxidation of pyridine salts in comparison with the free amine. The pyridine ring is very strong, and addition of acids to nitrogen can, of course, loosen its bonds with the carbon atoms and, perhaps, other bonds in the molecule as well.

1. The presence of hydrogen atoms attached to nitrogen. Hydrogen atoms attached to nitrogen are, in general, highly reactive. They are readily replaced by halogens and by various organic radicals, and are eliminated under the action of nitrous acid, etc.

\[ \mathrm{CH_3NH_2 + Cl_2 \rightarrow CH_3NHCl + HCl} \]

\[ \mathrm{CH_3NH_2 + ClCOCH_3 \rightarrow CH_3NHCOCH_3 + HCl} \]

\[ \mathrm{CH_3NH_2 + O{=}NOH \rightarrow CH_3OH + N_2 + H_2O} \]

\[ \mathrm{(CH_3)_2NH + HON{=}O \rightarrow (CH_3)_2N{-}N{=}O + H_2O.} \]

In tertiary amines, hydrogen attached to nitrogen appears when the amines pass into the state of hydroxides or salts. Reactions of this hydrogen include, for example, the process of synthesis of choline from ethylene oxide with trimethylamine and water:

\[ \mathrm{(CH_3)_3N + HOH \rightleftarrows (CH_3)_3N\!\begin{matrix}/\mathrm{H}\\[-2pt]\backslash\mathrm{OH}\end{matrix}} \]

\[ \begin{array}{c} \begin{matrix} \mathrm{CH_2} & \\ | & \backslash \\ \mathrm{CH_2} & \!>\mathrm{O} \end{matrix} + \begin{matrix} \mathrm{H} \\ \backslash \\ \mathrm{HO} \end{matrix} \mathrm{N(CH_3)_3} \;\longrightarrow\; \begin{matrix} \mathrm{CH_2{-}OH} \\ | \\ \mathrm{CH_2} \\ | \\ \mathrm{N{-}(CH_3)_3} \\ | \\ \mathrm{OH} \end{matrix} \end{array} \]

For some reactions of amine oxidation, the direct participation of aminic hydrogen has also been proved. Oxidation [7] of amines of the paraffin series apparently proceeds through radicals of the type \(\mathrm{RNH}\) and \(\mathrm{R_2N}\), and in the first case, perhaps, also \(\mathrm{R{-}N}\).

Earlier extensive studies of the oxidation of primary and secondary amines by potassium permanganate led to the establishment that the first reaction is the loss of hydrogen to the free radical, with subsequent conversion into an amine and an imine:

\[ \mathrm{RNH_2 \rightarrow RNH{-}} \]

\[ \begin{array}{c} \mathrm{R}\backslash \\ \mathrm{R_1}/ \end{array} \mathrm{NH} \;\longrightarrow\; \begin{array}{c} \mathrm{R}\backslash \\ \mathrm{R_1}/ \end{array} \mathrm{N{-},} \qquad \mathrm{2(CH_3)_2NH} \begin{matrix} \nearrow \mathrm{RCH_2NH_2} \\ \searrow \mathrm{R_1CH{=}NH} \end{matrix} \]

The products formed cannot be isolated, but on hydrolysis they are converted into ammonia and carbonyl compounds:

\[ \mathrm{R{-}CH{=}N + H_2O \rightarrow RC\!\begin{matrix} //\mathrm{O}\\[-2pt]\backslash\mathrm{H}\end{matrix} + NH_3,} \]

\[ \begin{array}{c} \mathrm{R}\backslash \\ \mathrm{R_1}/ \end{array} \mathrm{C{=}NH + H_2O \rightarrow} \begin{array}{c} \mathrm{R}\backslash \\ \mathrm{R_1}/ \end{array} \mathrm{C{=}O + NH_3,} \]

and further oxidation of the aldehyde, etc. [8].

With the hydrogen attached to nitrogen begins the action of certain oxidizing agents on amines of the benzene series. Oxidation of aniline, for example, with permonosulfuric acid gives nitrosobenzene:

\[ \mathrm{C_6H_5NH_2 + 2H_2SO_5 \rightarrow C_6H_5N{=}O + 2H_2SO_4 + H_2O} \tag{2} \]

Oxidation of aniline by chromic mixture to aniline black dye in quinone proceeds through phenylazote, phenylquinonediimine, and other products:

\[ \mathrm{C_6H_5NH_2 + O \longrightarrow C_6H_5N.} \tag{1} \]

\[ (\mathrm{C_6H_5})_2\mathrm{N} \rightleftarrows \mathrm{C_6H_5N{=}C_6H_4{=}NH} \quad \text{etc.} \]

Careful oxidation of diphenylamine by potassium permanganate leads to diphenylazote and tetraphenylhydrazine:

\[ (\mathrm{C_6H_5})_2\mathrm{NH} \longrightarrow (\mathrm{C_6H_5})_2\mathrm{N} \]

\[ 2(\mathrm{C_6H_5})_2\dot{\mathrm{N}} \rightleftarrows (\mathrm{C_6H_5})_2\mathrm{N} — \mathrm{N}(\mathrm{C_6H_5})_2 . \tag{9} \]

The presence or absence of hydrogen atoms at nitrogen should, of course, also affect the rate of action of oxidizing agents. The acceleration of the oxidation of pyridine by permanganate in the form of salts may, of course, also be due to the fact that in the salts hydrogen appears at the nitrogen atom:

\[ \begin{array}{ccc} \begin{array}{c} \mathrm{CH}\\[-2pt] /\!\!\backslash\\[-2pt] \mathrm{HC}\quad \mathrm{CH}\\[-2pt] | \quad\quad |\\[-2pt] \mathrm{HC}\quad \mathrm{CH}\\[-2pt] \backslash\!\!/\\[-2pt] \mathrm{N} \end{array} & +\ \mathrm{H_2SO_4}\ = & \begin{array}{c} \mathrm{CH}\\[-2pt] /\!\!\backslash\\[-2pt] \mathrm{HC}\quad \mathrm{CH}\\[-2pt] | \quad\quad |\\[-2pt] \mathrm{HC}\quad \mathrm{CH}\\[-2pt] \backslash\!\!/\\[-2pt] \mathrm{H{-}N{-}OSO_2OH} \end{array} \end{array} \]

The principal aim of the planned work was to clarify the influence on the rate of oxidation of amines and ammonium compounds of the presence of a hydrogen atom bound to nitrogen. To clarify this influence, amines were taken in which, in the form of salts with acids, there is a hydrogen atom and the corresponding tetrasubstituted salts (and in alkaline solution—bases).

The following compounds were investigated:

1. Triethylamine nitrate
   \[ (\mathrm{C_2H_5})_3\mathrm{N} \begin{matrix} /\,\mathrm{H}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

2. Methyl-triethylammonium nitrate
   \[ (\mathrm{C_2H_5})_3\mathrm{N} \begin{matrix} /\,\mathrm{CH_3}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

3. Dimethylaniline nitrate
   \[ \mathrm{C_6H_5N(CH_3)_2} \begin{matrix} /\,\mathrm{H}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

4. Trimethylphenylammonium nitrate
   \[ \mathrm{C_6H_5N(CH_3)_2} \begin{matrix} /\,\mathrm{CH_3}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

5. Diethylaniline nitrate
   \[ \mathrm{C_6H_5N(C_2H_5)_2} \begin{matrix} /\,\mathrm{H}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

6. Dimethyl-ethyl-phenylammonium nitrate
   \[ \mathrm{C_6H_5N(CH_3)_2} \begin{matrix} /\,\mathrm{C_2H_5}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

7. Pyridine nitrate
   \[ \mathrm{C_5H_5N} \begin{matrix} /\,\mathrm{H}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

8. Ethylpyridinium nitrate
   \[ \mathrm{C_5H_5N} \begin{matrix} /\,\mathrm{C_2H_5}\\[-2pt] \backslash\,\mathrm{NO_3} \end{matrix} . \]

9. Pyrrole, \(\mathrm{C_4H_4NH}\).

10. Aniline, \(\mathrm{C_6H_5NH_2}\).

11. Acetanilide,
    \[ \mathrm{C_6H_5NH{-}C(=O){-}CH_3}. \]

In addition to the amines and ammonium salts directly related to the solution of the main question of interest to us, several other compounds were also taken, the oxidation of which it was likewise of interest to determine. The rate of reaction with pyrrrole permanganate was determined in order to compare the stability of the pyridine and pyrrole rings, and that of acetanilide with aniline, in order to find out to what extent introduction of the acetyl group strengthens the aniline molecule.

The oxidation of these compounds was carried out with potassium permanganate at a temperature of \(18^\circ\text{C}\) in an aqueous medium (bidistillate), in the presence of caustic soda and sulfuric acid. The concentration of the substance under study and of the oxidizing agent was \(0.03\ \text{m/l}\). The unreacted \(KMnO_4\) was acidified with \(H(KI + H_2SO_4)\); the iodine liberated was titrated with hyposulfite of concentration \(0.1\ N\). The rate of reaction was determined from the oxidizing agent remaining after definite intervals of time.

Experimental Part

1. Preparation of amines and ammonium salts

Of the amines and salts selected for oxidation, aniline, pyrrole, and acetanilide were taken in ready-made form. The first two were purified by distillation, the last by crystallization from water. Ready-made salts were not available; therefore they had to be synthesized. The ammonium salts were obtained by addition of an alkyl halide with subsequent replacement of the halide by a nitro group. In preparing the salts, heating was carried out on a water bath, since in the salts at a temperature of \(180^\circ\) rearrangement occurs:

\[ \begin{array}{c} \text{[structural formula: phenyl-}N(CH_3)_3 I \rightarrow \text{methyl-substituted phenyl-}N(CH_3)_2HI \rightarrow \text{dimethyl-substituted phenyl-}NHCH_3HI \rightarrow \text{trimethyl-substituted phenyl-}NH_3I \end{array} \]

All the quaternary ammonium salts obtained were readily soluble in water; therefore recrystallization was carried out from water. The purity of the product was checked for the presence of halide (Beilstein reaction) and for the \(Ag\) ion (by precipitation). In obtaining the pyridine salts it was necessary to take into account their strong hygroscopicity and low decomposition temperature. Ethylpyridine was prepared in an ampoule, and the water bath was heated to no more than \(60^\circ\). Pure triethylphenylammonium could not be obtained under our conditions, owing to its rapid oxidizability. On recrystallization it quickly turned yellow.

2. Determination of the rate of reaction with potassium permanganate

When comparing the rates of oxidation of two compounds of the aliphatic series, we have:

\[ (C_2H_5)_2N \begin{matrix} \nearrow NO_3\\ \searrow H \end{matrix} \]

Table 2

\[ (\mathrm{C}_2\mathrm{H}_5)_3\mathrm{N} \begin{matrix} \nearrow \mathrm{NO}_3\\[-2mm] \searrow \mathrm{CH}_3 \end{matrix} \]

On comparing these two series of numbers it is seen that, for example, to attain an oxygen consumption of up to 40% required, in the first case, 25 minutes, whereas in the second case—520 hours.

Of the compounds of the aromatic series we have:

\[ \mathrm{C}_6\mathrm{H}_5\mathrm{N}(\mathrm{CH}_3)_2 \begin{matrix} \nearrow \mathrm{H}\\[-2mm] \searrow \mathrm{NO}_3 \end{matrix} \]

Table 2a

\[ \mathrm{C}_6\mathrm{H}_5\mathrm{N}(\mathrm{CH}_3)_3 \begin{matrix} \\[-1mm] |\\[-1mm] \mathrm{NO}_3 \end{matrix} \]

\[ \mathrm{C}_6\mathrm{H}_5\mathrm{N}(\mathrm{C}_2\mathrm{H}_5)_2 \begin{matrix} \nearrow \mathrm{H}\\[-2mm] \searrow \mathrm{NO}_3 \end{matrix} \]

\[ \mathrm{C}_6\mathrm{H}_5\mathrm{N}(\mathrm{CH}_3)_2 \begin{matrix} \nearrow \mathrm{NO}_3\\[-2mm] \searrow \mathrm{C}_2\mathrm{H}_5 \end{matrix} \]

In the benzene series an even more pronounced picture of the difference in the rate of oxidation is obtained.

\[ \mathrm{C_6H_5N} \begin{matrix} \diagup \mathrm{NO_3}\\[-2pt] \diagdown \mathrm{H} \end{matrix} \]

\[ \mathrm{C_6H_5N} \begin{matrix} \diagup \mathrm{C_2H_5}\\[-2pt] \diagdown \mathrm{NO_3} \end{matrix} \]

In the oxidation of pyrrole it was not possible to determine precisely the percentage of oxygen consumed. During oxidation pyrrole imparts a brown coloration to the solution and thereby obscures the change in the color of the indicator. Approximately it may be said that after only 5–6 seconds it is oxidized to 30% or more.

The study of the rate of oxidation of aniline and acetanilide was carried out in three media—neutral, acidic, and alkaline.

Table 4

Neutral medium

\[ \mathrm{C_6H_5NH_2} \]

\[ \mathrm{C_6H_5NH{-}C(=O){-}CH_3} \]

Alkaline medium

\[ \mathrm{C_6H_5NH_2} \]

\[ \mathrm{C_6H_5NH{-}C(=O){-}CH_3} \]

Acid reaction

\[ \mathrm{C_6H_5NH_2} \]

\[ \mathrm{C_6H_5NH{-}C(=O){-}CH_3} \]

Conclusions

1. On the main question of the significance of the presence of a hydrogen atom at the nitrogen, the results of the experiments carried out give a quite definite answer: ammonium salts of the limiting series and aromatic ammonium salts are always oxidized considerably faster if there is at least one hydrogen atom at the nitrogen.

Table 5

As these examples show, the difference in the rate of reaction with permanganate of trisubstituted or tetrasubstituted ammonium salts

is measured as being on the order of hundreds of times greater and more. This proves that the hydrogen attached to nitrogen is, in the present case, of very great importance and, one must suppose, it is with it that the oxidation reaction begins.

1. Pyridine salts with a hydrogen atom attached to nitrogen are also oxidized more readily than the tetrasubstituted salts, but the difference in the reaction rate here is not so great—only threefold—and is less than the difference in the rate of oxidation of pyridine in acidic and alkaline solutions. This fact suggests that, in the case of pyridine, loosening of the ring upon increase in the valence of nitrogen is of great importance.

2. Pyrrole is oxidized many times more readily than pyridine, which is fully consistent with the greater activity of pyrrole itself, and of the pyrrole ring in general, under the action of the most varied reagents.

3. Comparison of acetanilide with aniline shows that the introduction of an acetyl group makes the aniline molecule considerably more resistant toward oxidizing agents.

LITERATURE

1. Gattermann.—Practical Work in Organic Chemistry, p. 237.
2. Karrer, P.—Course of Organic Chemistry. Oxidation of aniline in nitrobenzene, vol. 1, p. 455. 1938.
3. Karrer, P.—Course of Organic Chemistry, p. 538. 1938.
4. Nikonova, L. S.—Proceedings of the Siberian Chemical-Technological Institute, No. 2. 1931.
5. Tronov, B. V., and Nikonova, L. S.—ZhRFKhO, issue 4. 1928.
6. Tronov, B. V.—Journal of Applied Chemistry, vol. 13, No. 7. 1940.
7. Shlenk, V., and Bergmann, E.—Organic Chemistry, vol. 1, p. 247. 1936.
8. Shlenk, V., and Bergmann, E.—Organic Chemistry, vol. 1, p. 248. 1936.