Ind. Eng. Chem. Res. 2005, 44, 5439-5446
5439
Chemistry of Organic Nitrates: Thermal Chemistry of Linear and Branched Organic Nitrates Manuel A. Francisco* and John Krylowski ExxonMobil Research and Engineering Company, 1545 Route 22 East, Clinton Township, Annandale, New Jersey 08801
The objective of our research was to measure accurate rate constants for the thermal, unimolecular decomposition of organic nitrates. Our research confirms that the rate-determining step is homolytic cleavage of the weak O-N bond to form alkoxy radical and NO2, but the rate constants reported in the past are incorrect. The alkoxy radical and NO2 engage in secondary reactions that ultimately generate stable products such as carbonyl and nitro compounds. Infrared spectroscopy (IR) and gas chromatography/mass spectrometry (GC/MS) were used to monitor the time dependent loss of organic nitrate and to characterize the products of the thermal reaction. Past research indicates that oxygen slows the rate of homolytic O-N bond cleavage to form radicals. Our research shows that the rates are the same in nitrogen as they are in air. The reaction in air produces R-β unsaturated ketones/aldehydes, which are not generated in a nitrogen atmosphere. The unsaturated ketone/aldehydes complicate the IR analysis, giving the appearance that the loss of organic nitrate slows down. Unhindered, linear organic nitrates have lower reaction rate constants than hindered organic nitrates. Activation energies were found to be lower for hindered organic nitrates. The steric strain present in the hindered organic nitrates may account for the weaker O-N bonds and faster thermal reaction rates. Reaction rates for thermal decomposition under nitrogen were found to decrease as the viscosity of the solvent increased. Introduction Organic nitrates have been known for well over a century.1 The chemistry of organic nitrates began to be a significant research focus as far back as the 1930s because of their importance as explosives. Comprehensive reviews are available2-4 on their use as explosives and as propellants.5,6 Since that time, the importance of organic nitrates as reactive species in other areas of science and technology has been investigated extensively. Organic nitrates have been used as ignition improvers in automotive fuels.7-10 They are also known to form during the catalytic reduction of NOx by hydrocarbons in automotive combustion engine exhaust gas.11-14 Organic nitrates are also used as vasodilators.15-17 They have been found in the volatiles of cooked, cured pork18 and are active components of air pollution.19 The infrared spectra, Raman, and X-ray diffraction of several organic nitrates have been reported20-23 along with calculations of lower electronic states24 and other physical properties such as dipole moments,25 latent heats of vaporization,26 mass spectrometry,27,28 and gas chromatographic separation.29 Synthesis of organic nitrates can be accomplished via several well-known reactions.30-36 The physiological and hazardous properties of organic nitrates have been welldocumented.37 Organic nitrates have also been useful in stabilizing trichloroethylene, which is used in cleaning iron and aluminum38 and as catalysts with phosphoric acid ester in the manufacture of ketene from acetic acid.39 The diverse chemical properties of organic nitrates can be attributed to the weak O-N bond that is the site * To whom correspondence should be addressed. Telephone: (908)730-2335. Fax: (908)730-3198. E-mail:
[email protected].
of thermal and chemical reactivity. Organic nitrates are similar to hydroperoxides. The energy of activation for thermal fragmentation and unimolecular, homolytic cleavage of the O-O bond in ethyl hydroperoxide to form two radicals (CH3CH2O* and *OH) is approximately 37 kcal/mol.40,41 The energy of activation for cleavage of the O-N bond of ethyl nitrate to form two radicals ranges from 36.0 to 41.2 kcal/mol depending on the study.42-44 Reviews are available.45,46 Most of the kinetic measurements in the literature to date have been conducted using IR to monitor the disappearance of the organic nitrate with time or the evolution of the products of thermal chemistry such as NO and NO2. These studies showed that the rate of thermal, unimolecular decomposition is slower in air than it is in nitrogen. Our research demonstrates that the rates in air and nitrogen are the same. Initial research on the thermal chemistry confirmed literature studies43-48 that the rate-determining step (RDS) in the thermolysis of organic nitrate is homolytic cleavage of the weak O-N bond to form alkoxy radical and NO2 (Figure 1). The initial RDS is a simple unimolecular reaction, but the mechanism shows how complicated the subsequent chemistry is. These complexities have manifested themselves in our measurements of the RDS and led us to the discovery that past measurements of rate constants for this step are not accurate. The alkoxy radical and NO2 can engage in secondary reactions with the solvent or other molecules that ultimately generate more stable products such as carbonyl compounds (aldehydes, ketones, and carboxylic acids) and nitro compounds. Lower molecular weight organic nitrates may also be generated. The mechanism may be more complex than illustrated. Only reactions
10.1021/ie049380d CCC: $30.25 © 2005 American Chemical Society Published on Web 06/15/2005
5440
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005
Figure 1. Mechanistic pathways of the thermal chemistry of organic nitrates shows that the rate-determining step (RDS) is the unimolecular thermal fragmentation of the organic nitrate to form two radicals.
Figure 2. IR specta of octadecyl nitrate ester (ODN) and EHN (neat) show the characteristic nitrate stretch at 1638 cm-1.
Figure 3. Concentration vs integrated IR absorbance of ODN and EHN in hexadecane is linear.
for which there is significant evidence have been shown. Infrared spectroscopy and GC/MS were used in this study to monitor the time-dependent loss of organic nitrate at various temperatures and to characterize the major products of the thermal reaction. Results and Discussion Model organic nitrates used were octadecylnitrate (ODN) and 2-ethylhexyl nitrate (EHN). The IR spectra of the ODN and EHN are shown in Figure 2. The critical IR absorption that is monitored with time is the absorption at 1638 cm-1, which is the nitrate stretching frequency. This can vary slightly, depending on the
organic nitrate structure (the absorbance is integrated from 1610 to 1660 cm-1). The integrated absorbance of this peak was found to be linear over the concentration range of 0-5 wt % (Figure 3) in hexadecane. The range between 0 and 1.0 wt % is most important, as this is where most of the kinetic measurements were made. The integrated absorbance of the carbonyl peak for EHN was also found to be linear over the same concentration range (Figure 3). The integrated absorbance of the organic nitrate peak decreases with time during thermolysis. The time-dependent loss of organic nitrate by IR followed a first-order rate law for all experiments involving ODN and EHN.
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005 5441 Table 1. Kinetic IR Data for ODN and EHN with Time at 170 °C under Nitrogen and Air in Hexadecane time 1660-10 (min) cm-1
ln a/ a-x
% 1760-1690 1564-1545 completeness cm-1 cm-1
time 1660-10 (min) cm-1
ln a/ a-x
% 1760-1690 1564-1545 completeness cm-1 cm-1
0 15 30 45 60 75 90 105 120 135
6.02 5.5 4.9 4.4 3.91 3.49 3.11 2.73 2.45 2.19
0.000 0.090 0.206 0.313 0.432 0.545 0.660 0.791 0.899 1.011
0.00 8.64 18.60 26.91 35.05 42.03 48.34 54.65 59.30 63.62
-10.74 -9.71 -8.43 -7.44 -6.58 -5.8 -5.09 -4.36 -3.85 -3.39
ODN (Nitrogen) 0.04 150 0.1 165 0.15 180 0.15 195 0.16 210 0.16 225 0.18 330 0.17 360 0.18 420 0.16
1.98 1.78 1.61 1.44 1.28 1.14 0.51 0.4 0.32
1.112 1.218 1.319 1.430 1.548 1.664 2.468 2.711 2.935
67.11 70.43 73.26 76.08 78.74 81.06 91.53 93.36 94.684385
0 15 30 45 60 75 90 105 120 135 150 165 180
5.41 4.63 3.79 3.15 2.53 2.16 1.79 1.54 1.26 1.05 0.89 0.74 0.61
0.00 0.16 0.36 0.54 0.76 0.92 1.11 1.26 1.46 1.64 1.81 1.98 2.18
0.00 14.42 29.94 41.74 53.16 60.09 66.88 71.51 76.78 80.55 83.59 86.26 88.69
-0.01 0.01 0.04 0.06 0.07 0.09 0.10 0.11 0.10 0.10 0.10 0.10 0.10
EHN (Nitrogen) 0.02 195 0.06 210 0.07 225 0.08 240 0.07 255 0.07 270 0.07 285 0.07 300 0.07 315 0.07 330 0.07 345 0.07 360 0.06
0.53 0.44 0.37 0.34 0.26 0.23 0.19 0.16 0.14 0.11 0.10 0.11
2.32 2.51 2.68 2.77 3.03 3.18 3.34 3.53 3.65 3.85 3.98 3.90
0 15 30 45 60
2.97 2.78 2.62 2.44 2.26
0.00 0.06 0.13 0.19 0.27
0.00 6.27 11.89 17.67 23.93
0.29 0.62 0.81 1.11 1.45
0.03 0.11 0.20 0.27 0.34
ODN (Air) 75 90 105 120
2.09 1.91 1.77 1.62
0 15 30 45 60 75 90 105 120
5.69 5.14 4.50 3.82 3.23 2.81 2.41 2.12 1.88
0.00 0.10 0.23 0.40 0.57 0.70 0.86 0.99 1.11
0.00 9.63 20.93 32.86 43.17 50.57 57.56 62.69 66.95
0.02 0.27 0.40 0.48 0.54 0.60 0.65 0.72 0.75
0.02 0.28 0.41 0.46 0.49 0.51 0.53 0.54 0.56
EHN (Air) 135 150 165 180 195 300 330 360 390
1.71 1.57 1.45 1.32 1.23 0.77 0.68 0.62 0.56
Examples of the kinetic IR data for ODN and EHN as a function of time in minutes are shown in Table 1. This data set is typical of all experiments conducted. It shows the disappearance of the organic nitrate absorption at 1660-1610 cm-1 and the appearance of the carbonyl (1760-1690 cm-1) and nitro absorptions (15641545 cm-1) as a function of time at 170 °C under nitrogen. The third column in the tables is the natural log (ln) of the initial concentration of the organic nitrate (a) divided by the initial concentration of the organic nitrate (a) minus the concentration (x) at a given time during the reaction. The concentration of the organic nitrate as a function of time in these tables were calculated from the graphs show in Figure 3. Some experiments were run to about 60% completion, others experiments were conducted up to >90% completion with the same results. Experiments with ODN showed that the average rate constant under nitrogen at 170 °C was 0.481 h-1 and the error at 95% confidence was 2.5% (Table 2). When the reaction was conducted in air the average rate was 0.299 h-1, which is significantly slower. The error at 95% confidence for the air experiments was 3.5%. The rates do not change when the concentration is halved to 0.5% ODN. One of the experiments at 1 wt % ODN was carried out in hexadecane that had been equilibrated with water. The reaction rate and the
-3.02 -2.65 -2.34 -2.03 -1.76 -1.51 -0.38 -0.19 -0.06
0.19 0.19 0.18 0.18 0.17 0.17 0.17 0.18 0.17
90.21 91.86 93.12 93.72 95.15 95.82 96.47 97.06 97.41 97.88 98.13 97.97
0.10 0.09 0.10 0.11 0.08 0.09 0.09 0.10 0.10 0.09 0.09 0.09
0.07 0.06 0.07 0.06 0.05 0.06 0.05 0.06 0.06 0.06 0.06 0.06
0.35 0.44 0.52 0.60
29.65 35.56 40.53 45.33
0.98 1.13 1.04 2.02
0.40 0.46 0.50 0.53
1.20 1.29 1.37 1.46 1.53 2.00 2.12 2.22 2.32
69.85 72.44 74.57 76.81 78.39 86.41 87.98 89.14 90.19
0.82 0.87 0.92 0.99 1.04 1.54 1.75 1.90 2.07
0.58 0.61 0.63 0.65 0.67 0.74 0.75 0.75 0.75
Table 2. Rates of Loss of ODN in Hexadecane at 170 °C Show the Effects of Nitrogen (N2) and Air on the Rate of Loss of ODN test
k (1/h) under N2
k (1/h) under air
1 (1% ODN) 2 (1% ODN) 3 (1% ODN) 4 (1% ODN) 5 (0.5% ODN) 6 (H2O present) average error 95% confidence
0.503 0.476 0.461 0.496 0.476 0.478 0.481 0.0122 (2.5%)
0.2793 0.298 0.298 0.304 0.297 0.320 0.299 0.0104 (3.5%)
product distribution were unchanged by the presence of water, indicating that hydrolysis of the organic nitrate was not occurring. Rate constants for EHN in nitrogen are greater than those of ODN by a factor of 1.5. This is may be attributed to the steric hindrance around the nitrate functionality in EHN, which weakens the O-N bond. This hypothesis is corroborated by energy of activation (Eact) measurements, which will be discussed later. The rate constants for EHN thermolysis under nitrogen at 170 °C average 0.722 h-1 over 12 separate experiments, and the error at 95% confidence is 0.0326 h-1 or 4.5% in nitrogen. The rates in air average 0.562 h-1 over 12 experiments, and the error at 95% confidence is 0.047 h-1 or 8.4%. Like ODN, the rate constant for EHN in
5442
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005
Figure 4. Partial IR spectra of nitrate functionality in ODN and EHN as a function of time at 170 °C in hexadecane show the effects of nitrogen (N2) and air on the 1632 cm-1 absorption.
Figure 5. Integrated absorbances at 1638 and 1632 cm-1 of ODN at 170 °C in hexadecane show the effects of nitrogen (N2) and air.
air is lower than it is in nitrogen. There were no changes in rate when EHN was reacted at half concentration (0.5 wt % in hexadecane at 170 °C under nitrogen) or when water was included. The effect of air on the rate is consistent with literature studies,43-48 using IR analysis, which have shown that oxygen slows the rate of homolytic O-N bond cleavage to form radicals. The explanation given in the literature is that oxygen oxidizes NO generated in the secondary reactions (Figure 1) converting it back to NO2. The concentration of NO2 increases the rate of the reverse of the RDS (NO2 reaction with alkoxy radical to regenerate the organic nitrate) also increases. Oxygen can also oxidize organic nitrite created in the secondary reactions to organic nitrate (Figure 1). Both reactions involving oxygen would have the effect of slowing the rate of loss of organic nitrate by IR analysis. IR Analysis. Close examination of the IR spectrum between 1610 and 1660 cm-1 suggests another explanation. Figure 4 (left) shows that when the thermolysis of the organic nitrate is conducted in a nitrogen atmosphere, the IR absorption at 1638 cm-1 is a single
absorption, which decreases steadily with reaction time. When the thermolysis is conducted in air, however, a new absorption steadily grows in at about 1632 cm-1. When the IR is integrated from 1660 to 1610 cm-1, as it was in past literature studies, both the 1638 cm-1 and the 1632 cm-1 absorbances are integrated. The result is the rate of organic nitrate thermolysis appears to slow in air but in fact it does not. This study shows that the rates are the same in air as they are in nitrogen. Figure 5 shows what happens to the integrated IR absorbances at 1638 and 1632 cm-1 as a function of thermolysis time at 170 °C in both nitrogen and air. This information was obtained by using a standard computer software package to deconvolute the IR data to obtain the rate of change data for each absorbance independent of the other. It is clear that some of the 1632 cm-1 absorption is forming even when the reaction is run under a nitrogen atmosphere. This is probably due to not excluding air rigorously from the reaction. The concentration of the organic nitrate (1632 cm-1 absorption) is insignificant as compared to the unknown (1638 cm-1) even toward the end of the
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005 5443 Table 3. Rates of ODN Decomposition Are the Same in Nitrogen and Air When the Contribution of the 1632 cm-1 Absorption Is Removed X% C18H37ONO2 in hexadecane at 170 °C 1% 1% 1% 1% 0.5% average error (95% confidence)
nitrogen 0.468 0.522 0.492 0.494 0.0368 (7.4%)
Table 4. Summary of GC/MS Analysis of EHN after 1 h at 170 °C under Air and Nitrogen
air 0.515 0.444 0.498 0.486 0.492 0.487 0.0238 (4.9%)
reaction. The air experiment is very different. The concentration of the organic nitrate (1632 cm-1 absorption) quickly becomes very significant relative to the unknown (1638 cm-1 absorption) and after 190 min the concentration of the organic nitrate (1632 cm-1 absorption) dominates. It is interesting that in both cases the 1632 cm-1 absorption increases to a certain concentration and then maintains that level or decreases slightly. Both cases suggest that the 1632 cm-1 product is created by reactions that require oxygen and that production continues to a point and then stops or at some point the 1632 cm-1 product is destroyed as quickly as it is made. The IR spectrum of EHN is very similar to that of ODN in the 1660-1610 cm-1 region under nitrogen and air. Potential Products. One hypothesis is that the 1632 cm-1 product is a lower molecular weight organic nitrate, which results from secondary reactions of the ODN in the presence of air (Figure 1, rearrangement of the peroxynitrite), but it is difficult to explain why the concentration of the lower molecular weight organic nitrate would stabilize later in the reaction (Figure 5). Organic nitrites (Figure 1) have a characteristic doublet of absorptions that occur in about the same region of the IR as organic nitrates.49,50 It is possible that one of the doublet absorptions appears at 1632 cm-1 and the other is hidden beneath the large nitrate absorption at 1638 cm-1. Organic nitrites, however, are not consistent with the known chemistry and the observations made in this report. The concentration of NO would be lower in air than in nitrogen, which would result in less nitrite. The mechanism in Figure 1 shows that when oxygen is present organic nitrites are quickly converted to organic nitrates. One would expect more organic nitrite formation in nitrogen than in air. Other potential candidates are olefins.49,50 Other possibilities are selected aromatics;49,50 conjugated CdN double bonds and NdN double bonds; organic amides, amines, and ammonium ions; N-nitro amines; pyridines; and tropolones, but none of these can be rationalized by the mechanism in Figure 1. The most likely explanation is that the 1632 cm-1 absorption is due to the formation of a conjugated carbonyl compound. The carbonyls generated in the secondary reactions are oxidized to R-β unsaturated carbonyls. This is consistent with the mechanism shown in Figure 1 and with the observed results in Table 3. Condensation and polymerization reactions of the unsaturated carbonyl compound would account for the steady-state concentration of this material later in the oxidation (Figure 5). The rate constants as a result of deconvoluting the IR spectra of some of the experiments are show in Table 3. The nitrogen and air rates are based only on the timedependent loss of the 1638 cm-1 absorption.
GC/MS analysis corroborates the rates found by IR for ODN and EHN under nitrogen. The GC/MS analysis clearly differentiates ODN and EHN from all the reaction products; therefore, the analysis confirms that the rates are not slowed in air and that the structure of the 1632 cm-1 absorption in the IR is due to R-β unsaturated ketones. The rates for EHN at 170 °C under nitrogen average 0.804 h-1 (error at 95% confidence is 0.05 h or 6.2%) and under air it is 0.728 h-1 (error at 95% confidence is 0.043 h or 5.9%) over a total of 10 experiments. Table 4 summarizes the GC/MS analysis of the products of the reaction of ethylhexyl nitrate (EHN) at 170 °C under nitrogen and air. The major products and their relative amounts at later times are similar, but the product distribution becomes more complex with minor products that could not be identified. Table 4 shows the compound on the left, the atomic mass units (AMU) of each compound type, and the relative amounts found when the reaction was run in air or nitrogen as shown in the middle of the table under the air and N2 columns. The relative amounts of each compound is the relative intensity as found in the GC/MS analysis in thousands of counts (K). Many products observed by IR were identified by GC/MS. Those are organic nitro compounds and carbonyl compounds. Compounds listed as “rearranged” are isomers. Unsaturated ethylhexyl aldehyde, the third molecule down in the table appears again as the 10th molecule in the table and is labeled “rearranged”. This simply means that the GC/MS analysis identified two molecules in the mixture with different retention times but similar MS data. The MS data each have the same molecular ion of 126 amu but different fragmentation patterns which result from the difference in structure between isomers. The difference in structure is also the cause of the different retention times.
5444
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005
Since the exact structure of these compounds is not known, model compounds could not be used to generate GC/MS response graphs as a function of compound concentration. The response factors for a given compound type should be comparable from the air to the N2 case. As a result the relative amounts listed in Table 3 are reasonably good quantitative measures of the relative amounts of each compound type in air versus N2. There were products detected by GC/MS that are difficult to identify by IR. These are the hydrocarbons, olefins, alcohols, and hydroperoxides. A major difference between the air and the N2 experiment is that there is more of each of the major products generated in air. One explanation of this is that, in air, the radicals that result from the RDS and secondary reactions can be terminated by reaction with O2. This is not possible under nitrogen, and the radicals tend to end up oligomerizing. The end result is the air reaction tends to produce relatively lower molecular weight highly oxygenated products that are volatile enough to be detected by GC/MS. The reaction under nitrogen tends to produce less of these products and more products that are relatively higher molecular weight, lower in oxygen content that are less volatile and more difficult to detect by GC/MS. The GC/MS results agree with IR analysis that organic nitro compounds are produced in higher concentration in the air reaction. One reason for this could be the fact that NO generated by reduction of NO2 can be reoxidized in the air reaction back to NO2 (Figure 1). This keeps the concentration of NO2 higher in the air reaction favoring the formation of more organic nitro compounds. Oxygen, in the air case, also drives the radical chain propagating reactions of oxidation. This leads to higher concentrations of carbon radicals than the nitrogen case which can couple with NO2 to form organic nitro compounds. Hydroperoxides were found in relatively high concentrations only when the reaction was conducted in air. There were no hydroperoxides detected in the nitrogen reaction. The most likely candidates for the 1632 cm-1 product are R-β unsaturated carbonyls. The IR absorbance for this class of compounds is known to appear in the 1632 cm-1 range and they have been identified by GC/MS as major products in the air reaction (Table 3). Significantly smaller amounts have been detected by GC/MS in the nitrogen reaction, but this is probably not significant enough to show up in the IR. Oxidative environments such as the EHN reaction in air provide conditions for the transformation of carbonyl compounds and olefins produced by thermal fragmentation of organic nitrates into unsaturated carbonyl compounds.51-54 Figure 6 shows an IR analysis of the products from the thermolysis of ODN at 170 °C. The RNO2 and the CdO in the legend refer to organic nitrate and carbonyl products that are formed as a function of time in both nitrogen atmosphere and air. The products are carbonyl compounds and organic nitro compound, which agrees with the GC/MS analysis in Table 4. More carbonyl and nitro compounds form in the presence of air than in the nitrogen atmosphere. This is consistent with the mechanism in Figure 1, which shows that oxygen is necessary for the formation of carbonyl and nitro compounds. Some are still formed in the nitrogen atmosphere most
Figure 6. Effects of nitrogen and air on the stable products of ODN thermal decomposition.
likely because air was not totally excluded. EHN gives very similar results. The Eact values for the RDS of a linear organic nitrate (i.e., dodecyl nitrate, DDN), 42.3 kcal/mol, and a branched organic nitrate (i.e., EHN), 39.7 kcal/mol, were derived as shown in Figure 7. These values were obtained by conducting reactions of DDN and EHN under the same conditions as previously used to obtain rates but at three different temperatures (see the Experimental Section).55 The Eact for EHN is significantly lower. This corroborates the rate data, which suggests that the O-N bond in branched organic nitrates such as EHN is weaker than it is in linear organic nitrates such as DDN and ODN. Average rates for the thermolysis of DDN are 0.447 h-1 under nitrogen and 0.303 h-1 under air. These values are within experimental error of the values found for ODN under nitrogen (0.481 h-1) and air (0.299 h-1). DDN was substituted for ODN at this point in the study to confirm that linear organic nitrates have similar rate constants.56 The effects of solvent are shown in Figure 8. The effect is due to the higher viscosity of mineral oil (S150N) and squalane as compared to hexadecane, and it is welldocumented for other solvents in the literature.43,48 The diffusion of the alkoxy and NO2 radicals (Figure 1) slows, as the solvent becomes more viscous. This cage effect causes the reverse of the RDS to become more significant, and the two radicals recombine to form the starting organic nitrate. The overall effect is a slowing of the rate of thermolysis. Conclusions Contrary to literature studies, the rates for thermal decomposition of organic nitrates are the same in air as they are under nitrogen. The rates for thermal decomposition of branched organic nitrates are higher than they are for linear nitrates. The bond energy of branched organic nitrates is lower than it is for linear nitrates, which is in agreement with the rate data. The rates of organic nitrates are slower in more viscous solvents. This agrees with previous literature studies. The 1632 cm-1 absorption generated in air is due to the formation of R-β unsaturated carbonyl compounds. Future work will focus on the reactions of organic nitrates with metals and other organic compounds and the chemistry leading to the formation of unsaturated carbonyls.
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005 5445
Figure 7. Eact values for DDN and EHN.
Figure 8. Effects of solvent on the rates of thermolysis of DDN and EHN.
Experimental Section The thermal reactions were conducted with low concentrations of organic nitrate in pure organic solvent with no other chemical species present. All model compounds, reagents, and solvent were obtained from Aldrich Chemical and were >95% pure. The reaction conditions were simplified to avoid complications due to potential secondary reactions. Thermolysis reactions were done in a 500 mL flask with 0.5 and 1 wt % ODN and EHN in hexadecane as the solvent (total reaction volume was 250 mL). Squalane and mineral oil were substituted for hexadecane in selected experiments with EHN to probe the effects solvent viscosity. The entire reaction mixture is about 20 g, and it was heated at 170 °C in a three-neck glass flask. The flask was equipped with a magnetic stir bar, two thermometers (one to record the temperature of the reaction mixture and the other to record the temperature of the hot oil bath), and a reflux condenser to condense and return any volatiles back into the reaction mixture. The stirring rate was adjusted to 400 rpm for every experiment and the heater adjusted to heat the oil bath and maintain the reaction mixture at a constant temperature. Temperature control was quite good and varied less than 1 deg. The condenser was cooled with a flow of chilled water from the laboratory-chilled water system. The atmosphere in the condenser and reaction flask was controlled by keeping the entire system closed except for the outlet at the top of the condenser, which was allowed to exit through an oil trap. A slight positive pressure of air or nitrogen was introduced through one of the three necks of the flask. The flow of air or nitrogen blankets the reaction mixture
and flows up the condenser and through the oil trap maintaining a slight positive pressure of the atmosphere of choice throughout the system. The reaction mixture was always purged with nitrogen to remove any dissolved air in the case of the nitrogen atmosphere experiments. One-half gram samples were removed periodically for IR analysis. The samples were loaded into KBr cells (0.1 mm path length), and full IR spectra (4000-500 cm-1) were recorded on a Mattson series FTIR model Rs2 infrared spectrometer (class IIa Laser Product). GC/MS analyses were conducted on the following instruments. Rate constants were determined using a Hewlett-Packard 5890 series II gas chromatograph/ mass spectrometer with a Supelco SPB capillary column (30 m × 0.25 mm i.d. × 0.25 µm film thickness, Catalog No. 2-4034, Column No. 12537-02B) to analyze aliquots of the reaction mixture taken at different times. The injection port temperature was set at 200 °C, the initial oven temperature was set at 70 °C and increased to 290 °C at a rate of 15 °C/min. The column effluent to the detector was split 100/1. Helium was the carrier gas at a flow rate of 0.8 mL/min (6.3 psig). Total run time was 19.67 min. The detector temperature was set at 300 °C. The ionization was accomplished by electron impact at 70 eV. Products were identified and characterized on a time-of-flight GC/MS (Leco Pegasus II) equipped with a HP6890 GC. The GC was equipped with an RTX-5 column (10 m × 0.18 mm i.d. and 20 µm film thickness). Injection port temperature started at 50 °C for 0.1 min and was increased at a rate of 700 °C/min to 300 °C (held for 2 min). The initial GC oven temperature was 0 °C and was increased at a rate of 50 °C/min to 300 °C (held for 3 min). EHN was purchased from Aldrich Chemical Co. Additional experiments to measure activation energies were conducted at 160, 180, and 190 °C. These experiments were identical to the experiments conducted at 170 °C under nitrogen. Acknowledgment We thank many of our colleagues for helpful discussions and advice in editing this report, Prof. Alan Katritzky of the University of Florida in Gainsville for the synthesis of DDN, and H. H. Carstensen of ExxonMobil for helping with the mathematical deconvolution of the IR spectra. Literature Cited (1) Kopp, E. Ann. 1847, 64, 320. (2) Svatopluk, Z. Thermochim. Acta 1997, 290, 199-217.
5446
Ind. Eng. Chem. Res., Vol. 44, No. 15, 2005
(3) Chen, J. K.; Brill, T. B. Thermochim. Acta 1991, 181, 7177. (4) Lurie, B. A.; Svetlov, B. S.; Chemyshov, A. N. 9th Symposium on Chemical Problems Connected with the Stability of Explosives; 1993; pp 119-156. (5) Wheeler, W. H.; Whittaker, H.; Pike, H. H. M. J. Inst. Fuel 1947, 20, 137-156/159. (6) Diplock, B. R.; Lofts, D. L.; Grimston, R. A. J. R. Aeronaut. Soc. 1953, 57 (505), 19-28. (7) Clothier, P. Q. E.; Aguda, B. D.; Moise, A.; Pritchard, H. O. Chem. Soc. Rev. 1993, 22, 101-108. (8) Aubert, M.; Clerget, P.; Duchene, R. C. R. Acad. Agric. Fr. 1935, 201, 879. (9) Ferretti, P. Ric. Sci. 1943, 14, 217; Chem. Abstr. 1949, 43, 6805. (10) Heinze, R.; Marde, M.; Veidt, M. Oel Kohle 1941, 37, 422; Chem. Abstr. 1942, 36, 4993. (11) Mukaiyama, T.; Hata, E.; Yamada, T. Chem. Lett. 1995, 505-506. (12) Hata, E, Yamada, T. Mukaiyama, T. Bull. Chem. Soc. Jpn. 1995, 68, 3629-3636. (13) Blower, C. J.; Smith, T. D. Zeolites 1993, 13, 394-398. (14) Walker, A. P. Catal. Today 1995, 26, 107-128. (15) Torday, L. L.; Jensen, S. J. K.; Csizmadia, I. G. Med. Chem. Res. 2003, 12 (3), 139-146. (16) Buckell, F.; Hartry, J. D.; Rajalingam, U.; Bennett, B. M.; Whitney, R. A.; Thatcher, G. R. J. Chem. Soc., Perkin Trans. 1994, 3, 401-403. (17) McCleane, G. U.S. Pat. Appl. Publ. 2002, 11 pp. Cont.-inpart of U. S. 6,221,915. US 2002035157 A1 2002032. (18) Mottram, D. S. J. Agric. Food Chem. 1984, 32 (3), 343345. (19) Roberts, M. Atmos. Environ. 1990, 24A (2), 243-287. (20) Carrington, R. A. Spectrochim. Acta 1960, 16, 1279-1293. (21) Lecomte, J.; Mathieu, J. P. J. Chim. Phys. 1942, 39, 57. (22) Chedin, J. J. Phys. Radium 1939, 10, 445-454. (23) Pauling, L.; Brockway, L. O. J. Am. Chem. Soc. 1937, 59, 13. (24) Harris, L. E. J. Chem. Phys. 1973, 58 (12), 5615-5626. (25) Cowley, E. G.; Partington, J. R. J. Chem. Soc. 1933, 12521254. (26) Gray, P. A.; Pratt, M. W. T. J. Chem. Soc. 1957, 21632168. (27) Fraser, R. T. M.; Paul, N. C. J. Chem. Soc. 1968, 659663. (28) Fraser, R. T. M.; Paul, N. C. J. Chem. Soc. B 1968, 14071410. (29) Roberts, J. M.; Fajer, R. W.; Springston, S. R. Anal.Chem. 1989, 61, 771-772. (30) Marans, N. S.; Zelinski, R. P. J. Am. Chem. Soc. 1950, 72, 5330-5331. (31) Nichols, P. L., Jr.; Magnusson, A. B.; Ingham, J. D. Atmos. Environ. 1953, 75, 4255-4258. (32) Boschan, R.; Merrow, R. T.; Van Dolah, R. W. Chem. Rev. 1955, 55, 485-510. (33) McKillop, A.; Ford, M. E. Tetrahedron 1974, 30, 24672475.
(34) Olah, G. A.; Narang, S. C.; Pearson, R. L.; Cupas, C. A. Synth. Commun. 1978, 6, 452-453. (35) Katritzky, A. R.; Marzorati, L. J. Org. Chem. 1980, 45, 2515-2516. (36) Boschan, R.; Merrow, R. T.; Van Dolah, R. W. Chem. Rev. 1955, 55, 485-506. (37) Von Oettingen, W. F. NIH Bull. 1946, No. 186. (38) Klabunde, W, U.S. Patent 2,435,312 (Feb. 3, 1948); Chem. Abstr. 1948, 42, 2916. (39) Ketene from acetic acid (Societe Rhodiaceta) 1948, FR930682 19480202 Patent language and authors unavailable. (40) Kirk, A. D.; Knox, J. H. Trans., Faraday Soc. 1960, 56, 1296. (41) Kirk, A. D.; Knox, J. H. Proc. Chem. Soc. 1959, 384. (b) Benson, S. W.; O’Neal, H. E. Kinetic Data on Gas-Phase Unimolecular Reactions; 1970. (42) Adams, G. K.; Bawn, C. E. K. Trans. Fraraday Soc. 1949, 45, 494. (43) Levy, J. B. J. Am. Chem. Soc. 1954, 76, 3254-3257, 37903793. (44) Pollard, F. H.; Marshall, H. S. B.; Pedler, A. E. Trans. Faraday Soc. 1956, 52, 59. (45) Boschan, R.; Merrow, R. T.; Van Dolah, R. W. Chem. Rev. 1955, 55, 485-510 (46) Lurie, B. A.; Svetlov, S. B.; Chemyshov, A. N. Proceedings of the 9th Symposium on Chemical Problems Connected with the Stability of Explosives; 1993; pp 119-156. (47) Phillips, L. Nature 1950, 165, 564; 1947, 160, 753-754. (48) Levy, J.; Adrian, F. J. J. Am. Chem. Soc. 1955, 77, 20152016. (49) Landgrebe, J. A.; Heath, D. C.; et al. Theory and Practice in the Organic Laboratory; Lexington, MA, 1973; pp 257-259. (50) Carrington, R. A. G. Spectrochim. Acta 1960, 16, 1270129. (51) Baldi, M.; Milella, F.; Ramis, G.; Sanchez, E. Vicente; Busca, Guido, Appl. Catal., A 1998, 166 (1), 75-88. (52) Yamada, O. Shigen Kankyo 1993, 2 (5), 513-19. (53) Havel, J. J. J. Am. Chem. Soc. 1974, 96 (2), 530-3. (54) Tuazon, E. C. Atkinson, R. Int. J. Chem. Kinet. 1990, 22, 1221-1236. (55) Note: three different temperatures were used to determine approximate values of Eact. Four or more different temperatures would give more rigorous values of Eact. The purpose of this was to see if the Eact values agreed, directionally with the rate constants. (56) Note: for a rigorous comparison of linear versus branched organic nitrates, N-octyl nitrate would be the ideal comparison to 2-ethylhexyl nitrate. DDN and ODN are not rigorous comparisons for 2-ethylhexyl nitrate.
Received for review July 15, 2004 Revised manuscript received April 25, 2005 Accepted May 10, 2005 IE049380D
