Ind. Eng. Chem. Res. 1991,30, 1485-1488 Knozinger, H. In The Chemistry of the Hydroxyl Group; Patai, S., Ed.; Interscience: New York, 1971; Part 2, pp 641-718. Lowry, T. H.; Richardson, K. S. Mechanism and Theory in Organic Chemistry; 3rd ed.; Harper and Row: New York, 1987. Mok, W. S. L.; Leeomboon, T. C.; Antal, Jr., M. J.; Richards, G. N. Mechanism of Formation of 5-(hydroxymethyl)-2-furaldehyde from D-Fruche and Sucrose. Carbohydr. Res. 1990,199,91-109. Morrison, R. T.; Boyd, R.N. Organic Chemistry, 3rd ed.; Allyn and Bacon: London, 1978. Narayan, R.; Antal, Jr., M. J. In Supercritical Fluid Science and Technology; Johnston, K. P., Penninger, J. M. L., Eds.; ACS Symposium Series 406; American Chemical Society: Washington, DC, 1989; pp 226-241. Narayan, R.;Antal, Jr., M. J. Influence of Pressure on the AcidCatalyzed Rate Constant for l-Propanol Dehydration in Supercritical Water. J . Am. Chem. SOC. 1990,112,1927-1931. Pearson, D. E.; Sawyer, J. S.; Cleveland, J. H. Phosphoric Acid Systems, 10, The Generation of Hydrocarbon Fuels from Waste and Fermentation Compounds. Znd. Eng. Chem. Prod. Res. Dev. 1980,19,245-250. Quist, A. S.; Marshall, W. L.; Jolley, H. R. Electrical Conductances
1485
of Aqueous Solutions at High Temperature and Pressure II. The Conductances and Ionization Constants of Sulfuric Acid-Water Solutions from 0 to 800' and at Pressurea up to 4000 Bare. J. Phys. Chem. 1965,69,2726-2735. Ramayya, S.; Brittain, A.; De Almeida, C.; Mok, W.; Antal, Jr., M. J. Acid-Catalyzed Dehydration of Alcohols in Supercritical Water. Fuel 1987,66, 1364-1371. Saunders, W. H.; Cockerill, A. F. Mechanisms of Elimination Reactions; Wiley: New York, 1973; pp 221-274. Townsend, S. H.; Abraham, M. A,; Huppert, G. L.; Klein, M. T.; Paspek, S. C. Znd. Eng. Chem. Res. 1988,27, 143. Whalley, E. Chemical Reactions in Solution Under High Pressure. Ber. Phys. Chem. 1966, 70,956968. Winter, 0.; Eng, M.-T. Hydrocarbon Process. 1976, Nov. Xu, X.; De Almeida, C.; Antal Jr., M. J. Mechanism and Kinetics of the Acid-Catalyzed Dehydration of Ethanol in Supercritical Water. J. Supercrit. Fluids 1990, 3, 228-232. Received for review July 3, 1990 Revised manuscript received January 14, 1991 Accepted January 23, 1991
Kinetics of Nitration of Anisole in Aqueous Nitric Acid David J. Belaon* Peterborough Regional College, Park Crescent, Peterborough, Cambridgeshire PE1 402,England
Values of pseudo-first-order rate constant and activation energy are reported for the nitration of anisole (methoxybenzene) in aqueous nitric acid in the concentration range 25.18-31.38 mol % at temperatures in the range 293-328 K. Correlation of rate constants with acidity function indicates that the reaction is diffusion controlled under these conditions. However, comparison of the relative rate with those of other aromatics suggests it may be on the borderline between complete diffusion control and encounter pair reaction control. The reaction may also be showing kinetics initially transitional between zero and first orders a t higher concentrations and temperatures. Activation energy decreases with increasing acid concentration. Introduction It has been established (Belson and Strachan, 1989)that the mechanism of aromatic nitration in aqueous nitric acid is the same as that in aqueous sulfuric acid, viz., H30+ + HN03 + 2H20 + NOz+
(1)
NO2++ ArH
(2)
encounter pair
encounter pair
benzenonium ion
fmt
products
(3)
In a large excess of nitric acid, the reactions of benzene, toluene, p-xylene, mesitylene, and certain quaternary ammonium ions are first order. For the more reactive pxylene and mesitylene, the slopes of plots correlating rate constant with acidity function (log kl' - log aHNOs versus -Ho) are close to 2.0, and this indicates a diffusion-controlled reaction, i.e. (2) is rate-determining, For the other compounds, which are less reactive, these slopes are higher, closer to 3. Here, the reaction of the species in the encounter pair (3) is rate-determining. Anieole (methoxybenzene) is another more reactive aromatic, but, in this case, kinetics transitional between zero and first order have been reported (Draper and Ridd, 1981). Under zero-order conditions, it is the rate of production of nitronium ion (1)that determines the rate of *Current address: Department of Chemistry and Physics, Nottingham Polytechnic, Clifton Lane, Nottingham N G l l 8NS, England.
reaction. Draper and Ridd give rate constant values for only two concentrations of nitric acid and at one temperature (298 K). No attempt was made to correlate rate constants with acidity function. The present work extends the range of HN03 concentrations used to enable this to be done and also reports data for a range of temperatures and values of activation energy. Experimental Section Materials. Concentrated nitric acid (-70% w/w) was of AnalaR grade and anisole and hydrazinium sulfate were of General Purpose Reagent grade. They were all supplied by British Drug Houses Ltd. and used without further treatment. The concentration of the concentrated nitric acid was obtained by a hydrometer method, and more dilute solutions were prepared by mixing together the requisite amounts of concentrated acid and water, as described elsewhere (Belson and Strachan, 1989). About 450 g of dilute solution was made up at a time, and to this 0.1 g of hydrazinium sulfate was added. Kinetic Method. A small amount (0.7 pL) of anisole was added by microlitre syringe to 20.0 cm3of nitric acid solution in a flask held in a water bath controlled to f0.1 "C.Some of the mixture was quickly poured into a l-cm cell, and this was inserted into the cell holder of a CE272 UV spectrophotometer (Cecil Instruments Ltd.). Water from the water bath circulated through the cell holder. The absorbance (A) of 340 nm was recorded at suitable intervals of time. The remainder of the mixture in the
O888-6885/91/ 2630-1485$O2-5010 63 1991 American Chemical Society
1486 Ind. Eng. Chem. Res., Vol. 30,No. 7,1991 Table I. Types of In (A. - A ) versus Time Plots at Various Temperatures (K) [HN08]/(mol 9%) 293.6 K 298.0 K 303.0 K 308.0 K 313.0 K 25.18 1 1 1 2 26.68 1 1 1 2 2 2 1 1 28.14 2 2 29.75 1 1 2 31.38 1 1 2 0.40 'I
0.4
318.0 K 2 2 2 2
-0.2-
-0.3-
4*ol -0.10
-0.1s
Figure 1. Integrated first-order rate equation plot for nitration of anisole type 1 (25.18mol % HNOs at 308 K).
flask was heated for 5 min in a boiling water bath and then cooled back to the experimental temperature. The absorbance (A,) was then measured. Results and Discussion Catalysis by Nitrous Acid. Previous investigations (Draper and Ridd, 1981) suggested the occurrence of a severe catalytic effect on nitration of anisole, ascribed to HN02, and hydrazinium sulfate was added to the nitric acid used, to eliminate it. In the present work, analysis showed that a typical diluted nitric acid solution contained only 1.2 mg dm-3 nitrous acid as NOz-. Little difference was observed in graphs of percentage reaction versus time for this solution and for one of the same HN03 concentration treated to remove nitrous acid (Belson, 1989). However, hydrazinium sulfate was added to all solutions to ensure comparability with the results of Draper and Ridd. Order of Reaction. The rate curves obtained were steep and showed appreciable linearity, at least initially. However, attempts to obtain values of kd and k{, using an equation appropriate to such transitional kinetics (Sheats and Strachan, 1978), were unsatisfactory. The linearity of plots was poor and values of k( were always lower than the initial rates obtained from the slopes of the rate curves. Guggenheim's method for firsborder kinetics also gave poor linearity of plots, so plots of In (A, - A) against time were used. Two types were obtained: (1) linear throughout the whole run, as exemplified by Figure 1; (2) linear, but with initial deviation, as exemplified by Figure 2.
328.0 K 2
1
-0.1-
-0.4-
323.0 K 2 2 2
B
Tlmo/mln
Figure 2. Integrated first-order rate equation plot for nitration of anisole type 2 (25.18mol % HN03 at 323 K).
Table I shows how the type of plot obtained depended on nitric acid concentration and temperature. It is clear that the reaction is basically first order under these conditions, but that initial deviation occurs at higher temperatures and HN03 concentrations. The steepness and initial linearity of the rate curves may be mainly an artifact of the high spectrophotometricabsorptivity of nitroanisole. The negative deviations of rate constant values from those of first-order kinetics evident early in runs giving type 2 plots can be explained in two different ways. (a) The reaction is initially transitional between zeroth and first orders. This occurs at higher temperatures and concentrations, as would be expected. (b) The reaction is first order throughout, but the solution has cooled in the time in which it is removed from the water bath, poured into the cell, and inserted into the spectrophotometer cell holder. Apparent temperature drops were estimated from values of initial rate constants, rate constants for the linear sections of type 2 plots, and values of activation energy (Table 111),with use of the Arrhenius equation (Belson, 1989). They were between 1.0 and 5.7 K,which might reasonably have been expected to occur during transfer of the sample to the spectrophotometer. However, the evidence that these temperature drops might be real was not very convincing. There was no correlation between the size of the temperature drop and the concentration of HN03 for a given nominal temperature, which would suggest merely random fluctuations in the time taken to transfer the sample. On the other hand, for a given solution, the higher the nominal temperature, the larger were the temperature drops, and so the times taken to warm up and the initial transients should be longer. In fact, the transients were observed to decrease with increasing temperature for all solutions. The possibility that, at higher temperatures and HNOBconcentrations, the order is less than 1 initially cannot be ruled out. The concentration of anisole, limited by the experimental method used, was only ca.0.32 X lo4
Ind. Eng. Chem. Res., Vol. 30,No. 7, 1991 1487 Table 11. Valuer of the Pwudo-First-Order Rate Constant [k,’) for the Nitration of Anisole 103 kils-’ [HNOaI/ (mol 56) log -Horn 293 K 298 K 303 K 308 K 313 K 25.18 -0.150 3.10 0.0906 0.166 0.268 0.382 26.68 -0.045 3.21 0.188 0.325 0.513 0.646 28.14 0.060 3.32 0.370 0.615 1.01 1.61 29.75 0.163 3.425 0.808 1.33 1.89 3.12 31.38 0.268 3.53 1.27 1.90 2.82 4.48 ~~
0 Thio work (OX I b l o 1) 0 Drrpor and Ridd.1081
-‘“1 -2.8-
./
318 K 0.726 1.23 2.53,2.44 4.48
323K 1.26 1.74 3.54
328K 3.47
Table 111. Variation of the Activation Energy, E , with Nitric Acid Concentration for Anisole [HNOS]/(mol%) 25.18 26.68 28.14 29.75 31.38 E,/(kJ mol-’) 82.0 75.2 73.3 67.4 65.1 Table IV. Relative Rates of Nitration other workers’ data at 298 K present work, HNOdaq) CHSN02- (CHB)& aromatic ha)” Odaa)” HNO.(aa)b 298 K 318 K benzene 1 1 1 1 1 toluene 25 20 24 22 12 anisole 175 76 28 p-xylene 130 114 200 256 89 mesitylene 400 350 380 483 136 ~
-3.0-
loo k’jloo c*o.
-5.2-3.4-
“Hoggettet al., 1969. *Hanson et al., 1976. -5.6-
Strachan, 1976). For a given concentration of nitric acid, activation energy decreases in the order EJbenzene) > EJtoluene) > EJanisole) > E,@-xylene) > EJmesitylene)
-3.8-
1
-4.0
1 3.0
3.1
3.2
3.3
3.4
3.5
3.6
3.7
mol dnr3. It seems that higher anisole concentrations,such as those used previously by other workers, viz., (2.4-206) X lO-9 mol dm4 are needed before transitional kinetics are clearly indicated.
Relative Rates of Nitration. An estimate of relative reactivity may be obtained by combining the kl’values at 298 K, for benzene in 36.85 mol % HN03 and for anisole in 29.74 mol % HN03. The value of kl’(anisole)/k{(benzene) is given in Table IV, together with data previously published for other aromatics (Belson and Strachan, 1989)and values obtained by other workers for aqueous nitric acid (Hanson et al., 1976) and for other media (Hoggett et al., 1969).
Rate Constant Values and Correlation with Acidity Function Values of k{ were obtained from the linear parts of the integrated rate equation plots, like those shown in Figures 1 and 2. They are given in Table 11, along with values of log amos at 298 K (Redlich et al., 1968)and Hoat 293 K (Rochester, 1970). The use of these data has been justified before (Belson and Strachan, 1989). A graph of log kl’298 - log aHNO ma was plotted against -HOm3and comprises Figure 3. t h e two previously obtained data (Draper and Ridd, 1981)are also plotted on this graph for comparison. The two sets of data show reasonable agreement, but Draper and Ridd obtained their HN03concentrations by titrimetric analysis and also corrected their values to allow for the zero-order start to the reaction. Errors arising from these sources have opposite effects on the rate constant value (Belson, 1989). The gradient of Figure 3 is 2.07. As previously demonstrated for other aromatics (Belson and Strachan, 19891, NO2+is the nitrating agent. A gradient close to 2 has been found for another reactive aromatic, indicating a diffusion-limited reaction. Activation Energy Values. These, obtained from the gradients of plots of log k{ against the reciprocal of temperature, are given in Table 111. Activation energy decreases as the acid concentration increases, as found for other aromatics in aqueous nitric acid (Belson and Strachan, 1989)and in aqueous sulfuric acid (Chapman and
Conclusion Table IV shows that the relative rates found in this work are in broad agreement with previous findings. Although anisole appears less reactive in aqueous nitric acid than might be expected, it seems, in common with p-xylene and mesitylene, to be nitrated at a rate determined by the rate of diffusion together of ArH and NOz+molecules to form the encounter pair. Under the same circumstances, this should differ only slightly between different aromatics. The larger the molecular radius, the lower the rate of diffusion. The ratio kl’(mesitylene)/kl’@-xylene) from Table IV is 1.89 at 298 K,and this could reasonably be explained by intrinsic differences in diffusion rate. However, the ratio kl’(mesitylene)/kl’(anisole)is 6.36 at 298 K,and this is too large to be explained in this way. Anisole seems to be on the borderline of reactivity between complete diffusion control and encounter pair reaction control under these conditions. Other aromatics might be expected to pass through such a transition stage under certain conditions. Extrapolation of the data used (Rochester,1970;M i c h et al., 1968)gives, for pure nitric acid (100mol %), -Ho = 6.30 and log OHNO, = 1.93. Extrapolation of Figure 3 gives values of 2.728 for log k{ - log aHNOS and 4.7 for log k,’ (all at 298 K) for nitration of anisole at -Ho= 6.30. Figure 4 shows the plot of Figure 3 extrapolated to this point, together with similar extrapolations, of plots already reported (Belson and Strachan, 19891,for other aromatics.
-Ho Figure 3. Plot of log k i - log am%at 298 K versus -Ho at 293 K for anisole.
1488 Ind. Eng. Chem. Res., Vol. 30, No. 7, 1991
tration for each below which nitration is no longer diffusion controlled and these plots in Figure 4 have gradients greater than 2.0.
45 1
I
morityleno
P-xylone anisole
31
log k‘,-
2
/
log
Acknowledgment I thank the Royal Society of Chemistry for Research Fund grants, Peterborough Regional College for the use of facilities, and Dr. A. N. Strachan of Loughborough University of Technology for advice and guidance. Nomenclature k,’ = observed zero-order rate constant, mol % s-l kl’ = observed (pseudo-)first-order rate constant, s-l A = absorbance A, = final absorbance log aHNOs = logarithm of nitric acid activity Ho= Hammett acidity function E, = activation energy Superscripts 293, 298 = temperature, where relevant, K Registry No. Benzene, 71-43-2;toluene, 108-88-3;anisole, 100-66-3;p-xylene, 106-42-3;mesitylene, 108-67-8.
Literature Cited
-8-0
-H,
3.5
4.0
4.5
5.0
5.5
8.0
0.5
I
I
I
I
I
I
I
7.0 I
Figure 4. Extrapolated graphs correlating rate constants with acidity function.
N.B. A, B, and C in Figure 4 are the quaternary ammonium ions: (A) C,&N+(CH& (B) C6H6CH2N+(CH3)3; (C)4-CH3C&4CH2N+(CH3)3.The data for A are actually experimental results obtained (Draper and Ridd, 1981) at concentrations up to and including 100 mol % HN03. Figure 4 appears to show that, before the acidity of pure nitric acid is reached, the rates of nitration of benzene and toluene will become greater than those of anisole, p-xylene, and mesitylene. This will not be so, though, if nitration of benzene and toluene eventually becomes diffusion controlled at a sufficiently high acidity and the gradients of their plots decrease toward 2.0. The evident curvature of the plot for the very low reactivity quaternary ammonium ion A, seen in Figure 4, indicate that the transitional behavior suggested above is being shown by the experimental data for this compound. The cases of anisole, particularly, and of p-xylene and mesitylene exemplify the opposite situation. There may be a nitric acid concen-
Belson, D. J. Aromatic Nitration Using Aqueous Nitric Acid. Ph.D. Thesis, Loughborough University of Technology, Loughborough, England, 1989. Belson, D. J.; Strachan, A. N. Aromatic Nitration in Aqueous Nitric Acid. J. Chem. Soc., Perkin Trans. 2 1989,1519. Chapman, J. W.;Strachan, A. N. Two Phase Nitration of Chlorobenzene in 79.8% Sulphuric Acid. In Industrial and Laboratory Nitrations; Albright, L. F., Hanson, C., Eds.; American Chemical Society: Washington, DC, 1976;pp 219-224. Draper, M. R.; Ridd, J. H. Nitration in Aqueous Nitric Acid. The Rate Profile and the Limiting Reaction Rates. J. Chem. Soc., Perkin Trans. 2 1981,94-99. Hanson, C.; Pratt, M. W. T.; Sohrabi, M. Some Aspects of Aromatic Nitration in Aqueous Systems. In Industrial and Laboratory Nitrations; Albright, L. F., Hanson, C., Eds.; American Chemical Society: Washington, DC, 1976; pp 225-242. Hoggett, J. G.; Moodie, R. B.; Schofield, K. Electrophilic Aromatic Substitution. 11. Nitration of Some Reactive Aromatic Compounds in Sulfolane and Nitromethane. J. Chem. SOC.B 1969, 1-11. Redlich, 0.; Gargrave, W. E.; Krostek, W. D. Thermodynamics of Solutions. XII. Consistency of Sparse Data. Activities of Nitric and Perchloric Acids. Ind. Eng. Chem. Fundam. 1968, 72 (2), 211-214. Rochester, C. H. The Hammett Acidity Function. In Acidity Functions; Academic Press: London, 1970. Sheats, G. F.; Strachan, A. N.; Rates and Activation Energies of Nitronium Ion Formation and Reaction in the Nitration of Toluene in 78% Sulphuric Acid. Can. J . Chem. 1978, 56 (9), 1280-1283. Received for review June 27, 1990 Accepted February 20, 1991
