L. S. Nikonova.

Rate of Oxidation of Amines by Potassium Permanganate under Various Conditions.

(Article 2.)¹

In the present work the oxidation of the following amines has been studied:
$\mathrm{CH_3NH_2}$, $(\mathrm{CH_3})_2\mathrm{NH}$, $(\mathrm{CH_3})_3\mathrm{N}$, $\mathrm{CH_3{-}CH_2{-}NH_2}$, $(\mathrm{CH_3{-}CH_2})_2\mathrm{NH}$,
$(\mathrm{CH_3{-}CH_2})_3\mathrm{N}$, $(\mathrm{CH_3{-}CH_2{-}CH_2})_2\mathrm{NH}$, $(\mathrm{CH_3{-}CH_2{-}CH_2})_3\mathrm{N}$,
$(\mathrm{CH_3{-}CH_2{-}CH_2{-}CH_2{-}CH_2})\mathrm{NH}$,

\[ \begin{array}{ccccc} & \mathrm{CH} & & & \\ \mathrm{CH} & | & \mathrm{C{-}CH_2{-}NH} & & \\ \mathrm{CH} & & \mathrm{CH} & & \\ & \mathrm{CH} & & & \end{array} \qquad \begin{array}{ccccc} & & \mathrm{CH{-}CH_3} & & \\ & \diagup & & & \\ \mathrm{CH_2} & | & & \mathrm{CH_2} & \\ \mathrm{CH_2} & | & & \mathrm{CH{-}NH_2} & \\ & & \mathrm{CH_2} & & \end{array} \]

Alkaloids: quinidine, piperine, coniine, atropine, quinine, brucine, caffeine, nicotine, morphine, codeine, cinchonine.

The experimental conditions were in general the same as in the preceding work. The amines and the oxidizing agent were taken in equimolecular quantities. The concentration of permanganate was somewhat higher (0.05 mole per liter). The temperature of the experiments was 18–20°.

The oxidation was carried out:

1) In a weakly acid medium, with the addition, per molecule of amine, of 1 mole of sulfuric acid (formation of the acid salt);

2) In an almost neutral medium, with the addition, per molecule of amine, of $1/2$ molecule of sulfuric acid (normal salt);

3) Without addition of acids or alkalis (the alkaline medium produced by the amine itself);

4) In a strongly alkaline solution, with addition of 1 mole of $\mathrm{NaOH}$.

The reaction products were not studied. According to the literature, the action of potassium permanganate on an amine leads to cleavage of the latter; nitrogen is liberated in the form of ammonia, and the hydrocarbon residues give aldehydes and ketones².

I measured the rate of reaction by titrating the oxidizing agent remaining after definite intervals of time.

Results of experiments:

Table 1. $\mathrm{CH_3NH_2 + KMnO_4 + NaOH}$

Table 2. $\mathrm{CH_3NH_2 + KMnO_4 + H_2O}$

Table 3. $\mathrm{CH_3NH_2 + KMnO_4 + 1\,M.\ H_2SO_4}$

Table 4. $\mathrm{CH_3NH_2 + KMnO_4 + \frac{1}{2}\,M.\ H_2SO_4}$

Table 5. $\mathrm{(CH_3)_2NH + KMnO_4 + NaOH}$

Table 6. $\mathrm{(CH_3)_2NH + KMnO_4 + H_2O}$

Table 7. $\mathrm{(CH_3)_2NH + KMnO_4 + 1\,M.\ H_2SO_4}$

Table 8. $\mathrm{(CH_3)_2NH + KMnO_4 + \frac{1}{2}\,M.\ H_2SO_4}$

$\mathrm{(CH_3)_3N + KMnO_4 + NaOH}$

When the reacting components are mixed, a green coloration appears at once, which very quickly changes to brown ($\mathrm{KMnO_4}$ is oxidized by approximately 50% per minute).

\[ (\mathrm{CH}_3)_3\mathrm{N} + \mathrm{KMnO}_4 + \mathrm{H}_2\mathrm{O} \]

Trimethylamine in aqueous solution is oxidized instantaneously; after 5 seconds the solution already assumes a brown coloration, skipping the green color.

Table 9. \((\mathrm{CH}_3)_3\mathrm{N} + \mathrm{KMnO}_4 + 1\,\mathrm{M}\ \mathrm{H}_2\mathrm{SO}_4\)

\[ (\mathrm{CH}_3)_3\mathrm{N} + \mathrm{KMnO}_4 + \tfrac{1}{2}\,\mathrm{M}\ \mathrm{H}_2\mathrm{SO}_4 \]

In 10 sec. the permanganate is reduced to a brown color, skipping green.

Table 10. \((\mathrm{CH}_3—\mathrm{CH}_2)\mathrm{NH}_2 + \mathrm{KMnO}_4 + \mathrm{NaOH}\)

Table 11. \((\mathrm{CH}_3—\mathrm{CH}_2)\mathrm{NH}_2 + \mathrm{KMnO}_4 + \mathrm{H}_2\mathrm{O}\)

Table 12. \((\mathrm{CH}_3—\mathrm{CH}_2)\mathrm{NH}_2 + \mathrm{KMnO}_4 + 1\,\mathrm{M}\ \mathrm{H}_2\mathrm{SO}_4\)

Table 13. \((\mathrm{CH}_3—\mathrm{CH}_2)\mathrm{NH}_2 + \mathrm{KMnO}_4 + \tfrac{1}{2}\,\mathrm{M}\ \mathrm{H}_2\mathrm{SO}_4\)

Table 14. \((\mathrm{CH}_3—\mathrm{CH}_2)_2\mathrm{NH} + \mathrm{KMnO}_4 + \mathrm{NaOH}\)

Table 15. \((\mathrm{CH}_3—\mathrm{CH}_2)_2\mathrm{NH} + \mathrm{KMnO}_4 + \mathrm{H}_2\mathrm{O}\)

Table 16. $(\mathrm{CH_3{-}CH_2})_2\mathrm{NH} + \mathrm{KMnO_4} + 1\,\mathrm{M}\ \mathrm{H_2SO_4}$

Table 17. $(\mathrm{CH_3{-}CH_2})_2\mathrm{NH} + \mathrm{KMnO_4} + \tfrac{1}{2}\,\mathrm{M}\ \mathrm{H_2SO_4}$

Triethylamine in the presence of alkali is oxidized very rapidly. After 15 sec. following addition of the oxidizing agent the solution already assumes a brown coloration. In aqueous solution, however, oxidation proceeds still more rapidly: only for one second does a green coloration of the solution appear, which instantaneously changes to brown.

Table 18. $(\mathrm{CH_3{-}CH_2})_3\mathrm{N} + \mathrm{KMnO_4} + 1\,\mathrm{M}\ \mathrm{H_2SO_4}$

Table 19. $(\mathrm{CH_3{-}CH_2})_3\mathrm{N} + \mathrm{KMnO_4} + \tfrac{1}{2}\,\mathrm{M}\ \mathrm{H_2SO_4}$

Table 20. Dipropylamine $+\ \mathrm{KMnO_4} + \mathrm{NaOH}$

Table 21. Dipropylamine $+\ \mathrm{KMnO_4} + \mathrm{H_2O}$

Table 22. Dipropylamine $+\ \mathrm{KMnO_4} + 1\,\mathrm{M}\ \mathrm{H_2SO_4}$

Table 23. Dipropylamine $+\ \mathrm{KMnO_4} + \tfrac{1}{2}\,\mathrm{M}\ \mathrm{H_2SO_4}$

Table 24. Tripropylamine + KMnO₄ + NaOH

Table 25. Tripropylamine + KMnO₄ + H₂O

Tripropylamine in the presence of sulfuric acid is oxidized more slowly. In 30 min. the decoloration of the permanganate reached 50%. The rate of oxidation with 1 M H₂SO₄ and with ½ M is almost the same. Intermediate titration proved extremely difficult, owing to the formation of lumps of a resinous substance strongly colored with iodine.

Table 26. Diisoamylamine + KMnO₄ + NaOH

(Diisoamylamine dissolves poorly in alkali).

In water \((\mathrm{C_5H_{11}})_2\mathrm{NH}\) also dissolves very poorly.

Complete decoloration of the permanganate to a brown color occurred in 8 minutes. In a sulfuric-acid medium the amine is oxidized more slowly. But owing to the precipitation of resinous lumps of substance, strongly colored with iodine, titration proved difficult. It was possible to titrate completely only one sample, taken after 8 min. 30 sec.; this gave 17.5%.

Table 27. Benzylamine + KMnO₄ + NaOH

Table 28. Benzylamine + KMnO₄ + H₂O

Table 29. Benzylamine + KMnO₄ + 1 M H₂SO₄

Table 30. Benzylamine + KMnO₄ + ½ M H₂SO₄

Table 31. Methylhexamethylenamine 1,3 + KMnO₄ + NaOH

Table 32. Methylhexamethylenamine 1,3 + KMnO₄ + H₂O

Table 33. Methylhexamethylenamine + KMnO₄ + 1 m. H₂SO₄

For the alkaloids I do not give numerical data, since they were oxidized too rapidly for the rate to be measured under the experimental conditions adopted.

The reaction-rate constants were not calculated. The following summary tables contain relative values of the rate, calculated from the intervals of time during which the same percentage of active oxygen is consumed.

The reaction rate taken as unity was: methylamine + KMnO₄ + 1/2 m. H₂SO₄.

Table 34. Relative rate of the reaction of amines with KMnO₄

Table 35. Effect of an acidic or alkaline medium on the rate of oxidation.

Table 36. Effect of the nature of the radical on the rate of oxidation under various conditions.

Table 37. Effect of the number of radicals on the rate of oxidation of amines under various conditions.

From the last summary tables the following conclusions may be drawn:

I. Effect of an acidic or alkaline medium (Table 35).

In all reliable experiments (with the exception of those in which the substance dissolves poorly and is titrated poorly), oxidation proceeds most readily in an aqueous medium without the addition of acids or alkalis. Amines were oxidized somewhat more slowly in the presence of NaOH, and oxidation proceeded considerably more slowly with the addition of one or half a molecule of sulfuric acid. An exception is trimethylamine, which with half a molecule of sulfuric acid is oxidized almost as rapidly as in an aqueous medium; this may be explained by the greater alkalinity of trimethylamine. Half-molecules of sulfuric acid should in general give neutral salts, but with strongly basic properties of the amine a somewhat alkaline reaction is possible.

II. Effect of the nature of the radical (Table 36).

The ratio of the oxidation rates of the various radicals depends on the reaction of the medium; in acidic and neutral solutions (in the form of the neutral salt, i.e., with \(1/2\) mole of \(H_2SO_4\)) the radicals are arranged in order of increasing rate of oxidation of the amine in the following series: methyl, ethyl, propyl, isoamyl, benzyl, methylcyclohexyl.

In an alkaline medium the series is, in general, the same, but with some exceptions. Amines with ethyl radicals in certain cases were oxidized more slowly than the corresponding methylamines.

III. Effect of the number of radicals (Table 37).

In an acidic medium, increasing the number of radicals in all cases accelerated oxidation. In alkaline solution exceptions were again observed, especially if the oxidizability per one radical is calculated; then, for example, for methylamines in the presence of \(aOH\) the following figures are obtained:

IV. All the alkaloids studied were oxidized (in the form of salts) considerably faster than simple amines. Those dissolving simply in water reacted extremely rapidly.

1. For Article 1, see B. V. Troonov and L. S. Nikonova, Zh. R. F.-Kh. O., vol. XI, pp. 541–549. ↩

2. O. Wallach; L. Claisen, Ber. 8, 1237 (1875); Haas, Rec. tr. ch., P.—B., 14, 166 (1895). ↩