Kinetics of the Nitration of Nitrobenzene with 100% Nitric Acid and Study of This System by Other Methods of Physicochemical Analysis

B. V. Tronov, Kh. A. Cheltsuk

All theories explaining the mechanism of the nitration process are divided by A. V. Topchiev into two main groups [1]. To one group he assigns all theories that proceed from the assumption that the nitrating agent is the molecule of nitric acid, which combines with the aromatic compound to form a primary complex, subsequently decomposing into the nitro compound and water. In this case it is assumed that the addition of nitric acid may occur either to the double bonds of the aromatic nucleus (Wieland), or, conversely, the aromatic nucleus adds to the molecule of nitric acid (Tronov, Nametkin and Zabrodina, Michael and Carlson).

To the other group may be assigned theories based on the assumption that, in the system aromatic compound—nitric acid, the nitrating agents are: a dimer of nitric acid, its complex with sulfuric acid, or cations arising from two molecules of nitric acid or in a mixture of nitric and sulfuric acids (Hantzsch, Lantz, Tronov, Usanovich, Titov, and others).

Many authors expressed the idea of nitration through the nitronium ion, \(NO_2^+\), but only in 1946 were these concepts supported experimentally and the existence of the nitronium ion finally confirmed [2—7].

The mechanism of addition of the nitronium ion to the aromatic ring and removal of the proton is at present represented in the following form [2, p. 226]:

\[ NO_2^+ + ArH \ \xrightleftharpoons[\ ]{\text{slow}}\ Ar^+\!\begin{matrix}/H\\ \backslash NO_2\end{matrix} \]

\[ Ar^+\!\begin{matrix}/H\\ \backslash NO_2\end{matrix} + A^- \xrightarrow{\text{fast}} ArNO_2 + HA \]

The first stage of this process should be further subdivided:

\[ \mathrm{NO_2^+ + ArH \longrightarrow Ar{-}H \longrightarrow Ar} \]
\[ \begin{array}{ccc} & \vdots & | \\ & \mathrm{NO_2} & \mathrm{NO_2 \ldots H^+} \end{array} \]

First, the nitronium cation adds as a complex to the carbon of the benzene ring, which should occur very rapidly. Then follows a more or less slow transition of the complex bond of carbon with nitrogen into a valence bond, while hydrogen, in the form of a proton, is transferred to the oxygen atom.

The present work was undertaken by us in order to study the process of nitration of nitrobenzene with absolute nitric acid.

It is known that nitrobenzene undergoes further nitration considerably more difficultly than benzene. The reason for this may be an increase in the stability of the molecule owing to a decrease in the electron density of all the links of the benzene ring under the influence of the nitro group. Such an influence must undoubtedly manifest itself. But a second reason is also possible, namely the binding of the nitrating reagent itself—whether \(\mathrm{HNO_3}\), its dimer, or the nitronium cation—by the nitro group into a complex, which should interfere with the action of the reagent on the hydrocarbon links.

To determine whether the formation of such complexes takes place, we decided to carry out a kinetic study of this system, as well as a study using a number of physicochemical methods of analysis. Viscosity, density, electrical conductivity, and refractive index were studied.

Experimental Part

For the study, chemically pure, absolutely anhydrous nitric acid of specific gravity 1.5477 at \(0^\circ\mathrm{C}\) was used. We obtained such acid by the method proposed by K. P. Mishchenko [8].

In a special apparatus under vacuum, a mixture of chemically pure nitric acid of specific gravity 1.4 with chemically pure sulfuric acid of specific gravity 1.84 was repeatedly distilled off (up to 4 times) on a water bath at a temperature of \(25—40^\circ\mathrm{C}\) at a residual pressure of \(25—30\) mm Hg.

Nitrobenzene was used after double distillation, \(d_4^{20}=1.2034\), and with boiling point \(210.9^\circ\) at 765 mm Hg.

The kinetic study was carried out at three temperatures—at \(0^\circ\), \(10^\circ\), and \(20^\circ\mathrm{C}\)—in the corresponding thermostats. Temperature regulation in the thermostats at \(0^\circ\) and \(10^\circ\mathrm{C}\) was carried out with an accuracy of \(\pm 1^\circ\), and in the thermostat at \(20^\circ\), with an accuracy of \(\pm 2^\circ\).

We investigated mixtures of nitrobenzene with nitric acid in the following ratios: per one molecule of nitrobenzene, 0.5; 1; 2; 3; 4 molecules of nitric acid.

In all cases, an equivalent amount of nitrobenzene in a glass flask with a ground-glass stopper was frozen at \(0^\circ\), and then, without removing the flask from the thermostat, an equivalent amount of nitric acid at \(0^\circ\mathrm{C}\) was added to the flask. Under the influence of the heat evolved, the nitrobenzene melted and a homogeneous mixture was obtained. The temperature during mixing rose only slightly, by \(1—2^\circ\), and the mixture then very quickly cooled again to \(0^\circ\). After this, the mixture was left in the thermostat for further study or transferred to thermostats at another temperature. Such mixing of the components prevented strong heating, which, as shown...

control experiments, which led to a considerable increase in the nitration rate at the initial moment.

We followed the rate of nitration from the consumption of nitric acid. For this purpose, samples were taken at definite time intervals and titrated with a solution of Ba(OH)\(_2\). The titration was carried out with phenolphthalein. A Ba(OH)\(_2\) solution was preferred to a NaOH solution, since with it it is easier to ensure the absence of CO\(_2\).

The viscosity, density, and electrical conductivity of the nitric acid—nitrobenzene system were studied at 0 and 10°; the refractive index at 10°C. The accuracy of temperature regulation was \(\pm 0.1^\circ\)C.

The viscosity was determined in an Ostwald viscometer of the closed type for volatile liquids. Mixing of nitrobenzene with nitric acid was carried out as described above for the kinetic study.

Discussion of Results

The data of the kinetic study are presented graphically in Figs. 1, 2, and 3.

Fig. 1. Rate of nitration at 0°C in the C\(_6\)H\(_5\)NO\(_2\)—HNO\(_3\) system at molar ratios: 2:1, 1:1, 1:2, 1:3, and 1:4.

Fig. 2. Rate of nitration at 10°C in the C\(_6\)H\(_5\)NO\(_2\)—HNO\(_3\) system at molar ratios: 2:1, 1:1, 1:2, 1:3, and 1:4.

Fig. 3. Rate of nitration at 20°C in the C\(_6\)H\(_5\)NO\(_2\)—HNO\(_3\) system at molar ratios: 2:1, 1:1, 1:2, 1:3, and 1:4.

The percentage of acid consumed was calculated, in systems with the ratios C\(_6\)H\(_5\)NO\(_2\):HNO\(_3\) = 2:1 and C\(_6\)H\(_5\)NO\(_2\):HNO\(_3\) = 1:1, relative to the total amount of acid in the sample taken; and in systems with the ratios 1:2, 1:3, 1:4, relative to one mole of acid, i.e., to the stoichiometric amount calculated for the formation of dinitrobenzene.

From analysis of the data it can be seen that, at small amounts of nitric acid, especially at 0.5 mole, and also at one mole of nitric acid per 1 mole of nitrobenzene, the nitration process proceeds very slowly.

The rate of nitration increases considerably in the syste—

... in the 1:2 system, and increases especially strongly in the 1:3 and 1:4 systems, although at first glance it would seem that in systems with an excess of nitrobenzene the acid should be consumed much more rapidly. The experimental data show the opposite.

B. V. Tronov and A. P. Grigor’eva [9] carried out a preliminary kinetic study of this same system, in which the nitric acid used was not absolute but fuming acid of specific gravity 1.52.

The results of this study also indicate a sharp increase in the rate of nitration when the number of moles of HNO₃ per mole of nitrobenzene is increased. B. V. Tronov suggested the formation of complexes of the compositions C₆H₅NO₂·HNO₃ and C₆H₅NO₂·2HNO₃.

Consequently, the considerable decrease in the rate of nitration of nitrobenzene can be explained mainly by the fact that the nitro group binds nitric acid into a complex.

From the fact that the nitration rate, at a ratio of two moles of nitric acid to one mole of nitrobenzene, increases markedly, it may be concluded that formation of a complex of composition C₆H₅NO₂·HONO₂ is possible; in this connection, the nitration process proceeds very slowly in the C₆H₅NO₂—HNO₃ systems with ratios of 2:1 and 1:1.

Since, with three and four moles of nitric acid, the nitration rate again increases considerably, it may be concluded that complexes of the following compositions are also formed: C₆H₅NO₂·2HNO₃ and C₆H₅NO₂·4HNO₃.

Complex formation must proceed through addition to the acidic atoms of the nitro group of the hydrogens of one or two molecules of nitric acid, the nitracidium cation H₂NO₃⁺, or, finally, nitronium:

$$ \mathrm{C_6H_5{-}N(=O){-}O \cdots HONO_2; \qquad C_6H_5{-}N(=O){-}O \cdots H_2NO_3^+;} $$

$$ \mathrm{C_6H_5{-}N(=O){-}O \cdots NO_2^+; \qquad C_6H_5{-}N(=O){-}O \cdots NO_2^+ \cdots NO_2^+.} $$

All these reactions hinder nitration of the benzene nucleus.

On the viscosity isotherm (Fig. 4, curve IV) at 0°C, and also at 10°C (curve III), there is a break at approximately 61 mole percent nitric acid and a small maximum within the range 80–85 mole percent. Such a course of the curves indicates that complex formation occurs when the components are mixed in this system. It may be considered that the composition of the complexes is as follows: C₆H₅NO₂·2HNO₃ and C₆H₅NO₂·4HNO₃.

The density isotherms (Fig. 4, curves I and II) show deviations from additivity. The refractive-index isotherm (Fig. 4, curve V) has a pronounced convexity directed away from the composition axis. All this confirms the formation of complexes.

The electrical-conductivity isotherms at 0 and 10° (Fig. 5, curves I, II), obtained by us, may be assigned to the case in which one component and the complex compounds formed conduct electric current poorly, while the other component is a good conductor (in our case, nitric acid). From the values of the specific electrical conductivity we calculated the relative temperature coefficient of electrical conductivity.

On the curve of the relative temperature coefficient of electrical conductivity \(z_0{}_{10}\) (Fig. 5, curve III) we have three maxima corresponding to the following contents of \(\mathrm{HNO_3}\), in mole percent—

Fig. 4. Density isotherms: I—10°, II—0°; viscosity isotherms: III—10°, IV—0°; V—refractive-index isotherm at 10°C.

Fig. 5. Electrical-conductivity isotherms: I—at 0°, II—10°C; III—curve of the relative temperature coefficient of electrical conductivity \(z_0{}_{-10}\)°C.

namely: 50; 66.7 and 80. This again indicates the formation of complex compounds of the composition: \(\mathrm{C_6H_5NO_2\cdot HNO_3}\); \(\mathrm{C_6H_5NO_2\cdot 2HNO_3}\); \(\mathrm{C_6H_5NO_2\cdot 4HNO_3}\).

Conclusions

1. The very large decrease in the rate of nitration of nitrobenzene in comparison with benzene is explained not only by the fact that, under the influence of the \(\mathrm{NO_2}\) group, the molecule becomes considerably more stable, but mainly by the fact that the nitrating reagent is bound by nitrobenzene into complex compounds.

2. From the data of the kinetic investigation, and also from data obtained by other methods of physicochemical analysis, it is possible to draw the conclusion that complexes of the following composition are formed: \(\mathrm{C_6H_5NO_2\cdot HNO_3}\); \(\mathrm{C_6H_5NO_2\cdot 2HNO_3}\); \(\mathrm{C_6H_5NO_2\cdot 4HNO_3}\).

Literature

1. Topchiev A. V. Nitration of Hydrocarbons and Other Organic Compounds. Publishing House of the Academy of Sciences of the USSR, 83, 1949.

2. Ingold K. K. Mechanism of Reactions and Structure of Organic Compounds. Foreign Literature Publishing House, 219, 1959.

3. Hantzsch A. Z. phys. Chem., 65, 41, 1908.

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