Kinetics and Mechanism of Direct Iodination of Aromatic Compounds in Acetic Acid Solution in the Presence of Nitric and Sulfuric Acids

A. M. SEDOV, A. N. NOVIKOV

(Submitted by the Methodological Seminar of the Chemical Technology Faculty)

Using the reaction of direct iodination of organic substances in the presence of a mixed sulfuric–nitric acid mixture, a considerable number of both iodo- and iodonitro derivatives of aromatic compounds have now been synthesized [1–3].

However, almost no attention has so far been paid to questions of the quantitative characterization of this process, i.e., determination of the reaction rate constant, partial orders, activation energy, and also the obtaining of information on the mechanism of the reaction under study. The present work was carried out for this purpose.

Methodology of the Kinetic Experiments

Benzene was chosen for the kinetic studies, since in its monoiodination side processes are absent.

Reagents: cryoscopic benzene; iodine of “pure” grade, twice sublimed; acids—glacial acetic acid, “chemically pure,” sulfuric acid \(d\ 1.83\), nitric acid \(d\ 1.4\). Solutions of iodine in acetic acid were prepared.

The component parts of the reaction mixture were kept in a thermostat until the required temperature was reached for 30 min, then rapidly combined. At the moment of mixing, a stopwatch was started. The reaction was carried out in vessels with ground-glass stoppers. The course of the process was followed by the change in iodine concentration in the mixture. At specified time intervals, 5 ml samples were taken and diluted fourfold with water. At this dilution the iodination reaction ceases completely. Iodine was extracted from the sample with two portions of \( \mathrm{CCl}_4 \), 11 ml each.

each, shaking vigorously for 2 min. The extracts were transferred to a 25-ml volumetric flask, CCl₄ was added to the mark, and the contents were mixed. The concentration of iodine in the CCl₄ solution was determined from the measured optical density using a calibration graph constructed from a series of standard solutions on a FEKN-57.

Thus, the experimental data of the kinetic studies represent values of the iodine concentration at different moments of time in the course of the reaction.

Determination of the rate constant \(K_j\) and the order of the reaction with respect to iodine

The Ostwald isolation method [4] was used. The following amounts of starting substances were taken: 50 ml of 0.020 mol/l acetic-acid solution of iodine, 0.02 mole of benzene (20-fold excess relative to the amount of iodine), 1.3 ml of HNO₃ (20-fold excess), and 2.6 ml of H₂SO₄ (45-fold excess). The temperature in the thermostat was 20°C. Samples of 5 ml were taken after 30 min. The results are presented in Fig. 1. All points lie on a straight line passing through zero. Good agreement is observed between the experimental results and the first-order equation [5].

Fig. 1. Graph for determining the order of the reaction with respect to iodine

Fig. 2. Graph for determining the half-conversion period of iodine (\(t_{1/2}\)). Curve 1—\([J_2]_0 = 0.020\) mol/l, 20°C; curve 2—\([J_2]_0 = 0.015\) mol/l, 20°C; curve 3—\([J_2]_0 = 0.010\) mol/l, 20°C; curve 4—\([J_2]_0 = 0.020\) mol/l, 25°C; curve 5—\([J_2]_0 = 0.020\) mol/l, 30°C; curve 6—\([J_2]_0 = 0.020\) mol/l, 35°C

\[ \ln \frac{a}{a-x} = K \cdot t . \tag{1} \]

\(K_j\) was determined from the half-conversion period \(t_{1/2}\) [6].

\[ K = \frac{0.693}{t_{1/2}} . \tag{2} \]

Different initial iodine concentrations were taken (0.020 mol/l, 0.015 mol/l, 0.010 mol/l). From the graph (Fig. 2) \(t_{1/2}\) was found and, using equation (2), \(K_j\) was calculated; it proved to be \(2.57 \cdot 10^{-3}\ \mathrm{min}^{-1}\) (\(t_{1/2} = 270\) min.

Determination of the reaction order with respect to benzene.

At \([I_2]_0=0.020\ \mathrm{mol/l}\), \([\mathrm{HNO}_3]=20[I_2]_0\), \([\mathrm{H_2SO_4}]=45[I_2]_0\), a 20–50-fold excess of benzene was taken. At all values of \([\mathrm{C_6H_6}]_0\), the half-conversion period of iodine is equal to 270 min; consequently, the reaction rate under these conditions does not depend on the concentration of benzene, and therefore a zero order of the reaction with respect to benzene was determined experimentally.

We also traced the effect of various amounts of nitric and sulfuric acids on the reaction-rate constant. With an increase in the amount of nitric acid, a decrease in \(K_1\) is observed (at \(20^\circ\mathrm{C}\) and a 60-fold excess of \(\mathrm{HNO_3}\), \(K_1=1.95\cdot 10^{-3}\ \mathrm{min}^{-1}\)); sulfuric acid considerably increases \(K_1\) (at \(20^\circ\mathrm{C}\) and an 80-fold excess of \(\mathrm{H_2SO_4}\), \(K_1=1.12\cdot 10^{-2}\ \mathrm{min}^{-1}\)).

Determination of the activation energy

The activation energy was calculated from the Arrhenius equation [4].

\[ \lg \frac{K_{T_2}}{K_{T_1}}=\frac{E}{4.575}\left(\frac{T_2-T_1}{T_2\cdot T_1}\right). \tag{3} \]

\(K_{T_1}=K_j\) at \(20^\circ\mathrm{C}\) is \(2.57\cdot 10^{-3}\ \mathrm{min}^{-1}=4.28\cdot 10^{-5}\ \mathrm{s}^{-1}\); \(K_{T_2}=K_j\) at \(30^\circ\mathrm{C}\), calculated from equation (2), has the value \(6.79\cdot 10^{-3}\ \mathrm{min}^{-1}=11.3\cdot 10^{-5}\ \mathrm{s}^{-1}\) (\(t_{1/2}=102\ \mathrm{min}\), Fig. 2). Substituting the values of \(K_j\) into equation (3), we found \(E=17.12\ \mathrm{kcal/mol}\). The activation energy was also determined graphically; for this purpose \(K_j\) was calculated at two more temperatures. \(K_j\) at \(25^\circ\mathrm{C}\) is \(4.22\cdot 10^{-3}\ \mathrm{min}^{-1}=7.04\cdot 10^{-5}\ \mathrm{s}^{-1}\) (\(t_{1/2}=164\ \mathrm{min}\), Fig. 2); \(K_j\) at \(35^\circ\mathrm{C}\) is \(10.81\cdot 10^{-3}\ \mathrm{min}^{-1}=18.05\cdot 10^{-5}\ \mathrm{s}^{-1}\) (\(t_{1/2}=64\ \mathrm{min}\), Fig. 2). Using the graph (Fig. 3), the activation energy was calculated by the formula [7]:

\[ E=4.57\cdot \tg\alpha\cdot \varepsilon, \]

where \(\varepsilon\) is the scale factor, equal to \((1:0.1\cdot 10^{-3}) : (1:1)=10^4\).

\[ E=4.57\cdot 0.3739\cdot 10^4=17.08\ \mathrm{kcal/mol}. \]

Fig. 3. Graphical determination of the activation energy.

Discussion of the results

On the basis of the results on the kinetics of benzene iodination, it may be concluded that the rate-limiting stage of the reaction studied is the interaction of iodine and the sulfuric–nitric acid mixture, since the rate constant depends on their amounts and does not depend on the concentration of benzene.

\[ \mathrm{I_2+HNO_3+H_2SO_4} \ \underset{}{\overset{\text{slow}}{\rightleftarrows}}\ \text{iodinating agent }(\mathrm{J^+}) \]

Probably, iodine polarized in the acetic-acid medium reacts with the nitronium cation, as may also be indicated by the obtained value of the activation energy. At the stage of electrophilic attack on the benzene ring

\[ \mathrm{C_6H_6 + I^+ \xrightarrow[\ ]{\text{fast}} C_6H_5I + H^+} \]

the formation of intermediate \(\pi\)- and \(\sigma\)-complexes is possible.

Conclusions

1. The kinetics of benzene iodination in the presence of a sulfuric–nitric acid mixture have been studied: the rate constant \(K_j\), the reaction order with respect to iodine and benzene, and the activation energy have been determined.

2. Considerations have been put forward concerning the probable mechanism of the reaction under study.

References

1. A. N. Novikov. ZhOKh, 24, 655, 1954; ZhOKh, 29, 58, 1959.

2. B. V. Tronov, A. N. Novikov. — “Izv. vuzov,” ser. khim. i khim. tekhnol., 3, 872, 1960.

3. V. T. Slyusarchuk, A. N. Novikov. ZhOrKh, 7, 1323, 1967.

4. Praktikum po fizicheskoi khimii [Practicum in Physical Chemistry], edited by S. V. Gorbachev. Izd. VSh., pp. 447, 454, 1963.

5. K. Laidler. Kinetics of Organic Reactions. “Mir,” 1966, p. 18.

6. K. B. Yatsimirskii. Kinetic Methods of Analysis, 2nd ed., Khimiya, Moscow, 1967, p. 17.

7. N. M. Emanuel, D. G. Knorre. Course of Chemical Kinetics. Izd. VSh., Moscow, 1962, p. 49.