CATALYTIC OXIDATION OF NITRIC OXIDE.

COMMUNICATION III.

Study of the Catalytic Activity of Activated Carbon.

N. P. Kurin and E. M. Novgorodova

The studies of Burdick¹), Szego and Guacci²), Boreskov and Shogam³), and Kurin and Blok⁴,⁵) demonstrated the possibility of increasing the rate of the reaction of oxidation of nitric oxide to dioxide by using sufficiently active solid catalysts. These authors investigated many substances as catalysts; however, only activated carbon from coconut shells and peach pits (Burdick), silica gel and aluminum oxide gel (Szego and Guacci), activated wood charcoal (Boreskov and Shogam), and silica gel, vanadium, and zinc-chromium catalysts, used respectively in the contact production of sulfuric acid and in the preparation of synthetic methanol (Kurin and Blok), were subjected to more detailed study. In these investigations, the kinetics of the process of oxidation of nitric oxide to dioxide were largely elucidated, and it was also established that the most active catalyst, when dry gas is passed through it, is activated carbon from coconut shells and peach pits. A drawback of such a catalyst, from the standpoint of its industrial use, is the fact that it is very vigorously poisoned by water vapor, and this deactivates it to the greatest degree, since industrial nitrous gases (in the production of nitric acid and, possibly, sulfuric acid by the nitrous process) always contain moisture. Phenomena of poisoning by water vapor are also characteristic of the other catalysts mentioned here. The only exception is the zinc-chromium catalyst, whose activity in the temperature interval 100–200°C increases with the introduction of water vapor into the gas being oxidized; however, the percentage of oxidation of nitric oxide observed under these conditions, both for dry gas and for gas containing water vapor, is low. In view of the foregoing, if one is to rely on the catalytic oxidation of nitric oxide, which is especially desirable for industrial practice, the problem of finding still more active catalysts that are not poisoned by water vapor remains highly urgent. In addition to its applied significance, the study of contact oxidation

NO to NO₂ on various catalysts is of theoretical interest connected with elucidating the mechanism of this process. Taking into account the increased catalytic activity of activated carbons, a detailed study of various grades of these substances is highly desirable, with a view to revealing the dependence of activity on the method of preparing the carbons. For this purpose we also undertook a study of the catalytic properties of activated carbon from a BN-type gas-mask carbon. This carbon, formed into small cylinders with \(d = 2\) mm and \(l = 3\text{--}5\) mm, possesses high mechanical strength, which is especially important for preventing large losses due to mechanical entrainment in industrial use.

The activity of the gas-mask carbon was studied by the dynamic method in an apparatus previously used by one of us (N. P. Kurin) for investigating the catalytic oxidation of NO to NO₂ in the presence of silica gel, vanadium, and zinc-chromium catalysts, and described in the first communication⁴). We sought chiefly to conduct the study of the process of catalytic oxidation of nitric oxide to dioxide under conditions as close as possible to isothermal, since only in this case can more definite conclusions be drawn regarding the kinetics of the given process. For this purpose the contact apparatus used were made in the form of a coil of thin-walled glass tubing with a diameter of 5 mm and a volume of 5.7 ml. The contact apparatus was filled with crushed activated carbon, the individual pieces having a diameter of 1.5–2 mm. To reduce to a minimum the oxidation of nitric oxide in the gas phase, before and after the contact apparatus the tubes supplying nitric oxide and air to the contact apparatus and removing the nitrous gases from it were made of capillary tubes. Nitric oxide was obtained by decomposition of sodium nitrite with sulfuric acid; after washing with alkali solution and concentrated sulfuric acid (to remove NO₂ and H₂O), the NO was mixed in the required ratio with air purified of CO₂ and dried. The dosing process for nitric oxide and air was carried out by means of carefully calibrated flowmeters. The oxidized gas after the contact apparatus was passed through concentrated sulfuric acid. By determining in the resulting nitrose, with a Lunge nitrometer, the total amount of nitrogen oxides and by titrating part of the nitrose with permanganate to determine the nitrous form of nitrogen, and also knowing the total amount of nitric oxide that had passed through the contact apparatus according to the flowmeter, it was possible to calculate the degree of oxidation of nitric oxide to NO₂. It should be noted that, when determining the nitrite form of nitrogen, in order to prevent large losses of nitrogen oxides due to volatilization, a certain amount of potassium permanganate was titrated with nitrose, and not vice versa. For this purpose a definite amount of 0.1 normal permanganate solution was placed in an Erlenmeyer flask and diluted to

Table 1

NO oxidation rate to NO₂ (in percent) in the presence of activated carbon, depending on air humidity, gas feed rate, initial concentration of oxide and temperature.

with a fivefold volume of warm water so that the temperature of the solution was \(-40^\circ\mathrm{C}\), and then, from burettes with a drawn-out long thin tip immersed in the permanganate solution, titration with nitrite was carried out until the permanganate was decolorized. The experiments were carried out with dry gas containing 1, 5, 3, 6, and 10% (by volume) nitric oxide—the remainder being air—at temperatures of 25, 50, 100, and \(150^\circ\mathrm{C}\), and at space velocities of 200, 400, 800, 1200, and 1600 volumes of the gas mixture \((0^\circ\) and \(760\ \mathrm{mm}\ \mathrm{Hg})\) per one bulk volume of catalyst per hour.

Experimental results and their discussion.

The results obtained are given in Table 1, and are also presented in Figures 1, 2, 3, 4, 5, 6, 7, and 8. In this table \(\alpha_{\mathrm{total}}\), as before, denotes the total degree of oxidation of NO to \(\mathrm{NO}_2\) in the presence of activated carbon, as found from the experiments. This quantity is determined by the rate of the reaction on the surface of the catalyst and by the rate of the reaction in the gas medium. \(\alpha_{\mathrm{gas}}\) denotes that degree of oxidation which would occur if the oxidation reaction of NO to \(\mathrm{NO}_2\) proceeded exclusively in the gas medium in the space not occupied by the catalyst, under conditions in which there is no catalytic acceleration of the reaction. The ratio \(\frac{\alpha_{\mathrm{total}}}{\alpha_{\mathrm{gas}}}\) shows how many times greater the total degree of oxidation in the presence of activated carbon is than the degree of oxidation of NO to \(\mathrm{NO}_2\) in the gas phase in the space not occupied by the catalyst. The degree of oxidation of NO to \(\mathrm{NO}_2\) in the gas phase in the presence of activated carbon was calculated from the well-known Bodenstein equation\(^6\). As is known, in order to calculate the degree of oxidation from this equation, in addition to the initial concentrations of NO and oxygen, it is also necessary to know the residence time of the gas in the reaction zone.

Determination of the residence time of the gas in a contact apparatus when using the dynamic method and a porous granular catalyst is associated with great difficulties, since the true free volume and the pore diameters of the grains of the contact substance are unknown. However, if in calculating the degree of oxidation in the gas medium one confines oneself to approximate values of the reaction time, then the latter can be found by dividing the free volume of the contact space not occupied by the catalyst by the second-by-second volume of gas at the temperature and pressure of the experiment. We determined the free volume for activated carbon by the same method as was described by Kurina and Blokh in communication 1, and this value proved to be equal to 65–70%. It should be noted, however, that the principal masses of gas pass through the space between the catalyst grains (if the catalyst grains were monolithic, then this free space would amount to only

Fig. 1. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to dioxide on space velocity and temperature in the presence of activated carbon. Initial concentration of NO in the feed gas: 1.5% (vol.).

Fig. 2. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to dioxide on space velocity and temperature. Initial concentration of NO in the feed gas: 3.0% (vol.).

Fig. 3. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to dioxide on space velocity and temperature. Initial concentration of NO in the feed gas: 6.0% (vol.).

Fig. 4. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to dioxide on space velocity and temperature. Initial concentration of NO in the feed gas: 10.0% (vol.).

Fig. 5. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to nitrogen dioxide on the initial concentration of NO in the feed gas and on the space velocity. Temperature 25°C.

Fig. 6. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to nitrogen dioxide on the initial concentration of NO in the feed gas and on the space velocity. Temperature 50°C.

Fig. 7. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to nitrogen dioxide on the initial concentration of NO in the feed gas and on the space velocity. Temperature 100°C.

Fig. 8. Dependence of the overall degree of oxidation ($\alpha_{\mathrm{overall}}$) of nitric oxide to nitrogen dioxide on the initial concentration of NO in the feed gas and on the space velocity. Temperature 150°C.

only 30–35%) and gas is forced through the internal pore volume only to a small extent. Gas exchange between the internal and external regions in an individual granule takes place mainly by diffusion, the rate of which, as is known, is low. Therefore, taking into account the difficulty of forcing gas through the free pore volume, as well as the slowness of diffusion, in calculating the residence time of the gas in the contact apparatus the free volume was taken to be 50%, i.e., the arithmetic mean of 30–35 and 65–70.

From the table and figures presented, it is evident that the activated carbon investigated by us, when used as a catalyst with dry gas, behaves in the same way as silica gel and the vanadium and zinc-chromium catalysts. Thus, at all the temperatures and initial NO concentrations investigated by us, as the space velocity of gas passage increases, the overall degree of oxidation of nitric oxide to dioxide decreases; moreover, this decrease still occurs considerably more slowly than the space velocity increases.

Thus, with a nitric oxide content in the initial gas of 1.5% (by volume) and a temperature of 25°C, increasing the space velocity from 200 to 1600, i.e., by a factor of 8, reduces the degree of oxidation only from 37.6 to 27.9%, i.e., by a factor of 1.35. A similar picture is observed at all other temperatures and initial concentrations of nitric oxide in the initial gas. This phenomenon, in which the decrease in the degree of oxidation lags behind the increase in space velocity, often observed in other heterogeneous catalytic processes⁷), is especially strongly expressed for the catalytic oxidation of nitric oxide in the presence of the grade of activated carbon investigated, as compared with the silica gel and zinc-chromium and vanadium catalysts studied by us earlier.

Turning to the determination of the effect of increasing temperature on the overall degree of oxidation, it should be noted that the activated carbon investigated by us behaves in the same way as silica gel and the other catalysts when operating with dry gas; that is, despite the presence of catalytic acceleration of the reaction, the overall degree of oxidation of nitric oxide to dioxide decreases with increasing temperature at all the space velocities and initial concentrations of nitric oxide in the initial gas investigated. Thus, in the presence of activated carbon as well, the temperature coefficient of nitric oxide oxidation, as for homogeneous oxidation, remains negative. It should be noted that the greatest decrease in the degree of oxidation occurs in the temperature range 25–100°C. With a further increase in temperature above 100°C, the decrease in the overall degree of oxidation (Figs. 1, 2, 3, and 4) occurs to a lesser extent. Further, from Figs. 5, 6, 7, and 8 it is seen that, with an increase in the initial concentration of nitric oxide at constant space velocity and temperature, the overall degree of oxidation ($\alpha_{\text{total}}$) increases within the range of NO concentrations investigated. Thus, for example, at a temperature of 25°C and a space velocity

of 200 (Fig. 5), with an increase in the initial concentration of nitric oxide from 1.5 to 10%, the overall degree of oxidation increases from 37.6 to 96%. This increase is especially pronounced in the region of low NO concentrations (up to 3%), where an almost direct dependence is observed between the concentration of NO in the initial gas and the observed degree of oxidation. With a further increase in the initial NO concentration in the initial gas, the degree of NO oxidation increases more slowly, apparently tending toward a certain limit, the absolute value of which is determined by the temperature, the space velocity, and the oxygen concentration (the latter quantity, as already indicated, remains practically constant in all experiments). The described character of the dependence of the overall degree of oxidation on the NO concentration in the initial gas is generally also observed at other temperatures, with the sole difference that at higher temperatures the curves \(\alpha_{\text{overall}} = f(\mathrm{NO})\) tend to straighten out. It is noteworthy that the points characteristic of 3% nitric oxide in the initial gas, at temperatures of 50, 100, and 150°C, do not fall on the curves \(\alpha_{\text{overall}} = f(\mathrm{NO})\). The reasons for this phenomenon have not yet been clarified. Finally, the table and graphs presented show that this grade of activated carbon accelerates the reaction of oxidation of NO to \(\mathrm{NO_2}\), i.e., it is a positive catalyst. Indeed, for all the conditions investigated by us, the ratio of the overall degree of oxidation \((\alpha_{\text{overall}})\) to the degree of oxidation in the gas phase, if no catalytic influence of the contact material were observed \((\alpha_{\text{gas}})\), is always greater than 1. The catalytic acceleration of the reaction of oxidation of nitric oxide to dioxide, conventionally expressed by the ratio

\[ \frac{\alpha_{\text{overall}}}{\alpha_{\text{gas}}} \]

at constant NO concentration, increases with increasing space velocity and with increasing temperature. Thus, for example, at an initial NO concentration of 1.5% and \(t = 150^\circ\mathrm{C}\), an increase in the space velocity from 200 to 1600 leads to an increase in the ratio

\[ \frac{\alpha_{\text{overall}}}{\alpha_{\text{gas}}} \]

respectively from 2.08 to 11.5 times; on the other hand, at the same concentration of nitric oxide and a space velocity of 1600, an increase in temperature from 25 to 150°C raises

\[ \frac{\alpha_{\text{overall}}}{\alpha_{\text{gas}}} \]

from 4.08 to 11.5 times. Similar relationships are also observed at all other temperatures, space velocities, and initial NO concentrations in the initial gas. At constant temperatures and gas space velocities, the ratio

\[ \frac{\alpha_{\text{overall}}}{\alpha_{\text{gas}}} \]

with an increase in the initial concentration of nitric oxide from 1.5 to 3.04 increases, and then, with a further increase of NO to 10%, decreases; i.e., the curves

\[ \left(\frac{\alpha_{\text{overall}}}{\alpha_{\text{gas}}}\right)_{T,\nu} = f(\mathrm{NO}) \]

have maxima.

Thus, for \(t = 150^\circ\mathrm{C}\) and \(V = 1600 \dfrac{\alpha_{\mathrm{total}}}{\alpha_{\mathrm{gas}}}\), at \(C_{\mathrm{NO,initial}} = 1.5;\ 3.6\) and 10%, respectively, they proved to be equal to 11.5; 15.0; 9.68 and 8.75. These same results indicate a greater acceleration of the reaction

\[ 2\mathrm{NO} + \mathrm{O}_2 = 2\mathrm{NO}_2 \]

at low (1.5 and 3.0%) rather than at high (6 and 10%) initial concentrations of nitric oxide.

Thus, from all of the above it is evident that the activity of the grade of activated carbon we studied is considerably less than the activity observed by Burdick for activated carbon prepared from coconut shells and peach pits (according to this investigator, the reaction was accelerated 500-fold) when thoroughly dried nitrous gas was passed through it, since even negligible traces of moisture led to a sharp decrease in the activity of these catalysts. The activated carbon studied by us was in operation for 2 months; however, its activity did not change substantially during this period. The catalyst also retained its mechanical properties.

To find a kinetic equation correctly expressing the relationship between the concentrations of the reacting substances and the time of catalytic oxidation of nitric oxide to nitrogen dioxide, we first tested the Bodenstein equation\(^6\), valid for the course of the reaction in the gas phase:

\[ K = \frac{1}{t}\frac{1}{(b-a)^2} \left\{ \frac{(b-a)\alpha}{a(1-\alpha)} + 2{,}303\log \frac{1-\alpha}{1-\frac{a}{b}\alpha} \right\} \tag{1} \]

where:
\(a\) — one-half of the concentration of nitric oxide,
\(b\) — the concentration of oxygen,
\(\alpha\) — degree of oxidation of nitric oxide,
\(t\) — reaction time.

To calculate the rate constants from equation (1), it is necessary to know the residence time of the gas in the reaction zone. As already indicated above, when using a granular porous catalyst, it is impossible to determine the reaction time exactly. Therefore, in equation (1), as in all subsequent equations, instead of \(t\) the reciprocal of the space velocity is substituted, corrected for the expansion of the gas with temperature, since \(t\) is proportional to

\[ \frac{1}{V\frac{T}{273}}, \]

and \(a\) and \(b\) were substituted in volume percent.

The results obtained from these calculations are presented in Table 2. From the table given, it is evident that the rate constants of the reaction of the catalytic oxidation of nitric oxide to dioxide in the presence of this grade of activated carbon, calculated

Table No. 2.

Rate constants of the reaction \(2\mathrm{NO} + \mathrm{O}_2 = 2\mathrm{NO}_2\) in the presence of anti-gas activated carbon, calculated by the Bodenstein equation

by equation 1, undergo a strong increase with increasing space velocity at all the temperatures and initial concentrations of nitric oxide investigated by us. This indicates the inapplicability of this equation to the catalytic oxidation of nitric oxide on this catalyst. Attempts were also made to apply the equation of Boreskov and Shogam—

\[ -\frac{dC_{\mathrm{NO}}}{dt} = K\frac{C_{\mathrm{NO}}^{1.5} C_{\mathrm{O}_2}}{C_{\frac{1}{2}\mathrm{NO}_2}}, \]

which at \(15^\circ\mathrm{C}\) satisfactorily expresses the results obtained for the catalytic oxidation of nitric oxide in the presence of the activated carbon studied by them. However, the calcul-

calculated from it, the rate constants for the catalyst we studied did not remain constant quantities, but changed markedly with increasing space velocity. This indicated that the character and mechanism of the catalytic oxidation of nitric oxide in the presence of a solid contact depend to a great extent on the nature and structure of the latter. Calculation of the rate constants from the equations proposed by T. Kurin and valid for the catalytic oxidation of NO to NO₂ on silica gel,

\[ \left(-\frac{dC_{\mathrm{NO}}}{dt} =K\frac{C_{\mathrm{NO}}^{1.5}\cdot C_{\mathrm{O2}}}{C_{\mathrm{NO2}}}\right) \]

and on a zinc-chromium catalyst

\[ \left(-\frac{dC_{\mathrm{NO}}}{dt} =K\cdot C_{\mathrm{NO}}^{1.5}\cdot C_{\mathrm{O2}}\right) \]

also did not lead to positive results, since the rate constants at constant temperature did not remain constant quantities, but were functions of the space velocity or, ultimately, functions of the reaction time.

In the further course of searching for a kinetic dependence for the catalytic oxidation of nitric oxide in the presence of countercurrent activated carbon, it was assumed that the following equation is valid for this case:

\[ -\frac{dP_{\mathrm{NO}}}{dt} = \frac{K P_{\mathrm{NO}}^{2} f(P_{\mathrm{O_2}})} {1+b_{1}\cdot P_{\mathrm{NO_2}}} \tag{2} \]

According to this equation, the rate of catalytic oxidation is directly proportional to the square of the partial pressure or concentration of nitric oxide; the reaction is inhibited by the reaction product, nitrogen dioxide. \(K\) is the reaction rate constant, and \(b_1\) is a coefficient associated with the degree of screening of the active catalytic surface by the reaction product. It should be noted that the rate of catalytic oxidation of NO to NO₂ must also be a function of the partial pressure or concentration of oxygen. However, if one takes into account the large excess of oxygen used in our experiments in comparison with nitric oxide, especially at its low initial concentrations, then it may be assumed with a high degree of probability that the rate of oxidation of NO to NO₂ will be determined by the concentration of NO and the concentration of the reaction product—NO₂, i.e. \(f(O_2)\) will be a constant quantity. An analogous case of the independence of the reaction rate from O₂ occurs in the contact oxidation of sulfurous gas to sulfuric anhydride on platinum.

Denoting \(K \cdot f(P_{O_2})\) by \(K_1\), we obtain:

\[ -\frac{dP_{\mathrm{NO}}}{dt}=\frac{K_1 P_{\mathrm{NO}}^{2}}{1+b_1 P_{\mathrm{NO}_2}} . \]

If we denote by \(A\) the initial amount of NO, and by \(x\) the amount of NO that has reacted by time \(t\), then it is convenient to write:

\[ \frac{dx}{dt}=\frac{K(A-x)^2}{1+bx}; \tag{3} \]

after integration and replacing natural logarithms by common logarithms, we obtain:

\[ K=\frac{1}{t}\, b \cdot 2.303 \log \frac{A}{A-X} -\frac{1}{t}(1+Ab)\frac{X}{A(1-X)} . \]

Replacing the unknown reaction time \(t\) obtained by us by the quantity proportional to it,

\[ \frac{1}{V\cdot \dfrac{T}{273}}, \]

and expressing \(x\) through the degree of oxidation \(\alpha\), \((x=A\alpha)\), we obtain:

\[ K_1 = V\frac{T}{273}\, b \cdot 2.303 \log \frac{1}{1-\alpha} - V\frac{T}{273}(1+Ab)\frac{\alpha}{A(1-\alpha)} . \tag{4} \]

This equation contains two constants, \(K_1\) and \(b\), and its applicability can be checked in the following way:

Introduce the following notation:

\[ Z= V\frac{T}{273}\, 2.303 \log \frac{1}{1-\alpha} \quad \text{and} \quad W= V\frac{T}{273}\frac{\alpha}{A(1-\alpha)}; \]

then

\[ W=\frac{b}{1+Ab}\,Z-\frac{K_1}{1+Ab}, \tag{5} \]

i.e., we obtain the equation of a straight line, in which \(\dfrac{b}{1+Ab}\) is the tangent of the angle of inclination of this line to the \(Z\)-axis, and the quantity \(-\dfrac{K_1}{1+Ab}\) is the intercept cut off by the line on the \(W\)-axis at \(Z=0\).

If all the assumptions made are correct, the quantities \(W\) and \(Z\), calculated from the experimental data, should be related by a rectilinear dependence. Therefore this simple method should serve as a reliable criterion for assessing the suitability of the given form of the kinetic equation for the reaction isotherm. The quantities \(W\) and \(Z\), calculated by the indicated method from the experimental data, are presented in columns 3 and 4 of Table (3) and are shown in Figs. 9, 10, 11, and 12. From these figures it is evident that, within the limits of experimental error, the quantities \(W\) and \(Z\) give straight lines. The constancy of the slope of the straight lines can also be established in the following way. From equation (5), the tangent of the angle of inclination of the straight line is

\[ \frac{b}{1+Ab}=\frac{W+\dfrac{K_1}{1+Ab}}{Z}. \]

Hence, by adding to the quantity \(W\) the intercept cut off on the ordinate axis \(\left(\dfrac{K_1}{1+Ab}\right)\), and dividing the algebraic sum by \(Z\), one should, at constant temperature and initial concentration of nitric oxide, obtain for all space velocities a constant value of the quantity \(\left(\dfrac{b}{1+Ab}\right)\). The results of these calculations are presented in columns 4 and 5 of Table 3. The constancy of the tangent of the angle of inclination of the straight line \(W=f(Z)\) is maintained quite satisfactorily. The quantities given in column 4 of Table 3, which are the intercepts cut off by the straight lines \(W=f(Z)\) on the ordinate axis at \(Z=0\), were determined graphically from Figs. 9, 10, 11, and 12. From the mean values of the tangents of the angles of inclination of the straight lines \(W=f(Z)\), given in column 6 of Table 3, and also using the relation according to which \(\operatorname{tg}\varphi=\dfrac{b}{1+Ab}\), the values of the coefficients \(b\) entering into kinetic equation (4) were calculated. The quantities \(b\) are given in column 7 of Table 3. From equation (4), into which the corresponding values of \(b\) and the observed degrees of oxidation of NO to \(\mathrm{NO_2}\) in the presence of activated carbon were substituted, the rate constants \(K_1\) of the reaction were calculated. The results of these calculations are summarized in column 8, and the mean value of \(K_1\) in column 9 of the same table. From this table it is evident that, at constant temperature and initial content of nitric oxide, the rate constants of the reaction \(2\mathrm{NO}+\mathrm{O_2}=2\mathrm{NO_2}\), calculated at different space velocities, remain constant within the limits of experimental error. The rectilinear character of the dependence between the quantities \(W\) and \(Z\), as well as the constancy of the rate constants calculated at different space velocities and at constant temperatures and initial concentrations, confirm the correctness of the assumptions we have made

Table 3

Continuation of Table 3.

Values marked with an asterisk (*) were discarded in calculating the mean value of the constant.

Fig. 9

Fig. 10

Fig. 11

Fig. 12

assumptions regarding the form of the kinetic dependence of the process of oxidation of nitric oxide to dioxide in the presence of activated carbon.

As Table 3 shows, the value of \(b\) changes both with increasing temperature and with increasing initial concentration of nitric oxide. Specifically, with increasing temperature the absolute value of \(b\) increases, although for an initial nitric oxide concentration of 10% in the temperature interval 50–150° the value changes little. An increase in the initial concentration of nitric oxide at all the temperatures studied leads to a decrease in the absolute value of \(b\); moreover, the greatest decrease is observed when the initial concentration of NO increases from 1.5 to 3.0%. The rate constants \(K_1\), calculated from equation (4), also change with increasing temperature and initial nitric oxide content. For temperatures of 25 and 50° the rate constants increase as the initial nitric oxide content increases. Thus, for example, at 25°C and initial nitric oxide concentrations of 1.5 and 10%, the rate constants are respectively 71.5 and 166.5. In the temperature region 100–150°C, the reaction rate constants, as the initial NO content increases, fluctuate around certain mean values. With an increase in temperature from 25 to 50°C at constant initial nitric oxide concentrations, the rate constants decrease; with a further increase in temperature to 150°C, the rate constants likewise fluctuate around certain mean values.

What has just been set forth indicates the very complex relationships that occur in the catalytic oxidation of nitric oxide to dioxide in the presence of activated carbon, if this process is considered over a wide interval of temperatures and initial nitric oxide concentrations. Taking into account the character of the temperature changes in the rate constants, it should be expected that the values of the apparent activation energy, as for the other catalysts studied, will be small values\(^{4}\). In this connection, the observed catalytic acceleration of the process of oxidation of NO to \(NO_2\) in the presence of the type of activated carbon investigated by us should be attributed not to a decrease in the activation energy, but chiefly to an increase in the concentration of the reacting substances on the active surface of the catalyst, caused by the presence of adsorption.

Thus, as before\(^{4}\), the observed acceleration of the reaction can be interpreted as the result of an increase in the number of collisions of active NO and \(O_2\) molecules, owing to their increased concentration on the surface of the catalyst. Taking into account the form of kinetic equation (3), in which the rate of oxidation of NO to \(NO_2\) in the presence of the activated carbon studied by us is directly proportional to the square of the concentration of nitric oxide, it may be assumed that the oxidation process proceeds

as in the gas phase, as a result of a triple collision of two nitrogen molecules and one oxygen molecule. However, for this it is necessary additionally to assume that the rate of oxidation of NO to NO\(_2\) is directly proportional to the oxygen concentration to the first power, since from our experiments it is not possible to reveal the influence of the oxygen concentration on the oxidation rate, owing to the large excesses of O\(_2\) used in comparison with NO. Some confirmation of the correctness of the observation made should be seen in the fact that, according to the studies of Boreskov and Shogam \(^{2}\), the rate of oxidation of NO to NO\(_2\) in the presence of the grade of activated carbon studied by them is directly proportional to the oxygen concentration to the first power.

Kurin and Blok came to the same conclusion in studying the catalytic oxidation of NO to NO\(_2\) in the presence of chromium and vanadium catalysts.

In considering the mechanism of the catalytic oxidation of NO to NO\(_2\) in the presence of the grade of activated carbon studied by us, as well as on other catalysts studied \(^{4}\), three cases may occur. The oxidation reaction of NO to NO\(_2\) proceeds as a result of:

1) a triple collision of two molecules of nitric oxide and one molecule of oxygen in the adsorption layer;

2) the impact, from the gas medium, of an oxygen molecule upon two NO molecules in double collision in the adsorption layer; and

3) an increase in the concentration, in the adsorption layer, of N\(_2\)O\(_2\) molecules, which, according to Bodenstein, should be regarded as an intermediate compound in the homogeneous oxidation of nitric oxide in the gas phase. According to Bodenstein, the concentration of N\(_2\)O\(_2\) in the gas medium is determined by the rapidly established equilibrium

\[ 2\mathrm{NO} \rightleftarrows \mathrm{N}_2\mathrm{O}_2. \]

The N\(_2\)O\(_2\) formed in this reaction then reacts with oxygen at a measurable rate according to the reaction

\[ \mathrm{N}_2\mathrm{O}_2 + \mathrm{O}_2 \to 2\mathrm{NO}_2, \]

according to which the total rate of oxidation of NO to NO\(_2\) will be directly proportional to the concentration of N\(_2\)O\(_2\) and to the oxygen concentration.

It is difficult to say which of the indicated cases actually occurs; it is possible that all are realized simultaneously, and each of them is consistent with the kinetic equation (3) found by us, according to which the apparent order with respect to nitric oxide for the catalytic oxidation of NO to NO\(_2\) in the presence of this grade of activated carbon is equal to 2, while the overall apparent order of the reaction, taking into account the influence of oxygen, is apparently equal to 3.

Further, the reaction

\[ 2\mathrm{NO} + \mathrm{O}_2 = 2\mathrm{NO}_2 \]

in the presence of activated carbon is inhibited by the reaction product NO\(_2\), which can be explained by a decrease in the free active surface of the catalyst as a result of its shielding by adsorbed NO\(_2\). The dependence of the quantity \(b\), entering into equation (4), on temperature and on the initial concentration of nitric oxide is caused—

is explained by a temperature-dependent change in the adsorption of $\mathrm{NO_2}$ on the surface of activated carbon and, apparently, is associated with the presence of energetically nonequivalent active centers.

Conclusions

1. Using the dynamic method, the activity of activated carbon from BN gas-mask charcoal for the catalytic oxidation of $\mathrm{NO}$ to $\mathrm{NO_2}$ was investigated at space velocities from 200 to 1600 and at temperatures from 25 to 150°C. The dry gas mixture fed to oxidation contained from 1.5 to 10% (by volume) nitrogen oxide, the remainder being air.

2. This grade of activated carbon proved to be a positive catalyst for accelerating the reaction $2\mathrm{NO} + \mathrm{O_2} = 2\mathrm{NO_2}$.

3. In the temperature range studied, with increasing temperature, at all the space velocities investigated, the overall degree of oxidation of $\mathrm{NO}$ to $\mathrm{NO_2}$ decreases.

4. For all the isotherms investigated, the overall degree of oxidation decreases with increasing space velocity of the gas. However, the decrease in the degree of oxidation occurs considerably more slowly than the space velocity increases. This leads to the greatest acceleration of the reaction, as compared with oxidation in the gas phase, precisely at high space velocities.

5. The overall process of catalytic oxidation of nitric oxide to dioxide under the investigated conditions, with a considerable excess of oxygen compared with the amount of $\mathrm{NO}$ in the presence of this grade of activated carbon, is expressed by the following kinetic equation:

\[ -\frac{d C_{\mathrm{NO}}}{dt} = K \frac{C_{\mathrm{NO}}^{2}}{1 + b C_{\mathrm{NO_2}}}. \]

It is assumed that the reaction of oxidation of $\mathrm{NO}$ to $\mathrm{NO_2}$ on the catalyst, as in the gas phase, is a trimolecular reaction, and that its observed acceleration is caused by an increase in the concentration of $\mathrm{NO}$ (or $\mathrm{N_2O_2}$) and $\mathrm{O_2}$ in the adsorption layer. In addition, the process of catalytic oxidation of $\mathrm{NO}$ to $\mathrm{NO_2}$ is inhibited by the $\mathrm{NO_2}$ formed as a result of the reaction.

Literature

1. Burdick, J. Am. Chem. Soc. 44, 244, 1922.
2. Szego, Gazz. Chim. Ital. 60, 212, 1930.
   Szego and Guacci, Gazz. Chim. Ital. 61, 333, 1931.
3. Boreskov and Shogam.—Zh. Fiz. Khim. 8, 306, 1936.
4. Kurin and Blok.—Zh. Prikl. Khim. 11, no. 6, 734, 1938.
5. Kurin and Blok.—Zh. Prikl. Khim. 11, no. 5, 750, 1939.
6. Bodenstein, Z. Elektrochem. 24, 183, 1918; Z. Angew. Chem. 31, 145, 1918; Z. Phys. Chem. 100, 87, 1922.
7. Kurin.—Zh. Khim. Prom. 7, 475, 1937.