Energy & Fuels 1988,2,729-734
729
Peroxide Formation in Jet Fuels George E. Fodor,* David W. Naegeli, and Karen B. Kohl Belvoir Fuels and Lubricants Research Facility, Southwest Research Institute, 6220 Culebra Road, San Antonio, Texas 78284 Received March 21, 1988. Revised Manuscript Received August 8, 1988 The rates of peroxide formation in six model jet fuels were measured at various temperatures ranging from 43 to 120 "C with oxygen partial pressures ranging from approximately 10 to 1140 kPa. One of the fuels exhibited an increase in the rate of peroxide formation after alumina treatment, and three of the fuels showed induction periods. The results agreed with a kinetic model of the autoxidation process in that the peroxide concentration increased as the square of the stress duration. The rate of peroxide formation did not depend on the oxygen partial pressure. Arrhenius correlations of global rate constants determined from peroxide concentration time histories in accordance with the kinetic model showed that a single autoxidation mechanism explains the results obtained in the 43-120 "C temperature range. The results of this work encourage the development of a test method that predicts rate of peroxide formation at ambient conditions from data that may be obtained from more timely experiments at elevated temperatures.
Introduction From 1962 to 1983,the United States Navy and some commercial air carriers experienced fuel system failures in jet aircraft flying in the western Pacific. Shertzer,' Hazlett, et al.? and Love, et a13assigned these difficulties to peroxides in the fuel. They found that peroxides cause significant deterioration of neoprene, nitrile rubber, and Buna-N diaphragms and O-rings used in the fuel pumps of jet engines. To avoid future problems, a program was initiated to study the kinetics of peroxide formation and ultimately to develop a timely method of predicting the potential peroxide content of jet fuels.'* It is well-known that peroxides form in fuels by an oxidation process that is relatively slow at room t e m p e r a t ~ r e . ~At ? ~ higher temperatures, the rate of fuel oxidation is dramatically increased. However, it is not certain that the reaction mechanism responsible for peroxide formation remains the same as that at ambient temperature^.^^' Several test methods have been developed to determine the oxidative stability of fuels, e.g., ASTM D 2274;they are carried out at elevated temperatures to reduce the test duration to an acceptable level. An accelerated high-temperature test is also desired for the determination of potential peroxide formation in jet fuels. Since the objective was to provide a basis for a practical test method, not exceeding 48 h, the foremost goal was to ~
(1)Shertzer,R. H. Final Report No. NAPC-433; Naval Air Propulsion
Center: Trenton, NJ, 1978. (2) Hazlett, R. N.; Hall, J. M.; Nowack, C. J.; Craig, L. C. Proceedings of the Conference on Long-Term Stabilities of Liquid hcekr; Israel Institute of Petroleum and Energy: Tel Aviv, Israel, 1983; No. B132. (3) Love, B. E.; Hatchett, K. A.; Peat, A. E. "Fuel-Related Problems in Engine Fuel Systems". SAE Trans. 1967, 75,441-463. (4) Fodor, G. E.; Naegeli, D. W.; Kohl, K. B.; Cuellar,J. P., Jr. Interim Report BFLRF No. 199, AD A163590; Belvoir Fuels and Lubricants Research Facility, Southwest Research Institute: San Antonio, TX,June 1985. (5) Fodor, G. E.; Naegeli, D. W. Proceedings of the 2nd Internationul Conference on Long-Term Storage Stabilities of Liquid Fuels; Southwest Research Institute: San Antonio, TX, 1986; pp 632-645. (6) Fodor, G. E.; Naegeli, D. W.; Kohl, K. B.; Cuellar, J. P., Jr., Interim Report BFLRF No. 243, AD A189293, Belvoir Fuels and Lubricants Resaarch Facility, Southwest Research Institute: San Antonio, TX,June 1987. (7) Watkins, J. M., Jr.; Mushrush, G. W.; Hazlett, R. N. Prepr. Pap.-Am. Chem. SOC.,Diu. Fuel Chem. 1987, 32(1), 513-521. (8) Walling, C. Free Radicals in Solution; Wiley: New York, 1957; Chapter 9. (9) Hine, J. Physical Organic Chemistry; McGraw-Hill: New York, 1956.
determine if the mechanism of peroxide formation at elevated temperatures was the same as that at ambient temperature. Assuming that the mechanism for peroxide formation does not change over a limited temperature range (e.g., 0-150 "C),it is theoretically possible to predict ambient temperature behavior from a global Arrhenius rate expression determined by making two or more rate measurements at higher temperatures. In the present study, the kinetics of peroxide formation in six kerosenes were examined over the temperature range 43-120 "C. Experiments in the 80-120 "C range were performed in a pure oxygen atmosphere in a stirred pressure reactor. In longer term experiments at 65 and 43 "C,bottle storage was used with fuels exposed to air.
Experimental Section To establish base-line data on the long-term stability,a modified version of the ASTM D 4625 method of bottle storage at 43 "C was used.I0 In the modified procedure, 300 mL of each fuel sample was purged with "synthetic" air (21% 0,and 79% N,) a t 300 mL/min until the fuel became saturated with oxygen, as determined by gas chromatography. Then the fuel samples were stored a t 43 "C in sealed 500-mL amber borosilicate bottles. This modification of the published procedure was necessary to detect and correct for possible depletion of oxygen in the bottled fuels. Such oxygen depletion was of special concern in the case of the more reactive fuels, where it could have had a significant effect on the rate of peroxide formation. After scheduled aging periods, one bottle of each fuel was retrieved for analysis of peroxides (ASTM D 3703),gum (ASTM D 381),water (ASTM D 1744),and acid number (ASTM D 664). The oxygen contents in both the liquid and vapor phases were determined by gas chromatography. If the oxygen concentration in the vapor phase dropped below 10.0 vol %, the remaining bottles of the same fuel were again aerated. The procedure described by Hall" was used for bottle storage experiments at 65 "C. This protocol was basically the same as that used in the 43 "C experiments, except that, in the 65 "C experiments, the samples were aerated by opening the bottles to room air for about 15 min weekly to alleviate oxygen depletion. Accelerated oxidative stressing at temperatures above 65 "C was carried out in replicate by using two nominally identical 600-mL, 316 stainless-steel pressure reactors. Temperature was regulated to k0.5 "C, and pressure was continuously monitored. (10) Standard Test Method for Distillate Fuel Storage Stabilityat 43 (110 OF), ASTM D 4625-86. (11)Hall, J. M. Memorandum to Participants in the Third CRC Cooperative Test Program on Hydroperoxide Potential of Jet Fuels, April 7, 1986. O C
0887-0624/88/2502-0729$01.50/00 1988 American Chemical Society
Fodor et al.
730 Energy & Fuels, Vol. 2, No. 6,1988
fuel 1 2 3 4 5 6
Table I. Model Fuels description straight-run, additive-free, salt-dried, clay-treated kerosene hydrocracked kerosene, alumina treated hydrocracked kerosene fuel 3, alumina-treated hydrofined kerosene hydrocracked kerosene
The reactors were each charged with 300 mL of fuel and purged with ultrahigh-purity oxygen (99.99%);they were then pressurized with pure oxygen and heated to the desired teat temperature. The fuel was stirred at 150 rpm to ensure rapid dissolution of oxygen and to prevent temperature gradients. The test temperature was reached in less than 30 min, and the clock was set to zero. Very little, if any, peroxide was formed in the time required to heat the fuel to the test temperature. In preliminary experiments, the oxidized fuel samples were analyzed for peroxides, gums, water, and acid number as was done with the bottle storage samples. These analysea were performed to provide information germane to the oxidation mechanism. However, in the majority of experiments, only peroxide content was measured, and an aliquot sampling procedure was adopted so that several samples could be obtained from a single-batch experiment. The f i t aliquot of sample was withdrawn from the reactor when the test temperature was reached. Sampling was then continued at convenient time intervals to measure the buildup of peroxide concentration. The test fuels selected for this study are described in Table I. Fuel 1 was a straight-run, salt-dried, clay-treated, and additive-free kerosene; it was chosen to serve as a pristine fuel of high oxidative stability. Four hydrocracked kerosenes were selected as potentially unstable fuels. Fuel 2, which contained about 400 ppm of peroxides when it was received, was percolated through alumina before use. The alumina treatment removed the peroxides and possibly other polar compounds. Another hydrocracked kerosene, fuel 3, was received with negligible peroxide content so it was not altered before use in the experiments. However, one purpose of the study was to examine the effect of alumina treatment on the formation of peroxides. To study this aspect of the problem, fuel 4 was prepared by alumina treating a sample of fuel 3. After the program was under way and considerable data had been obtained on the formation of peroxides in fuels 1-4, the investigation was extended to include fuels 5 and 6. These two fuels, received from the Naval Reaearch Laboratory, were part of the fuel matrix used in the third Coordinating Research Council (CRC) cooperative tegt program on hydroperoxide potential of jet fuels.ll The fuels were purged with argon and stored in sealed containers at 5 OC. All the test fuels were claimed by the suppliers to be free of added antioxidant type additives. Fuel 3 appeared to contain an antioxidant, but infrared analysis of the polar fraction of the fuels did not detect any substituted phenols or phenylenediamine-type compounds in excess of 5 ppm, the detectability limit of the procedure. The basic goal of this study was to determine if it is possible to predict the slow formation of peroxides in jet fuels at ambient conditions from experimental data of the relatively fast oxidation at higher temperatures. Rates of peroxide formation were measured over the temperature range 43-120 OC, and the partial pressures of oxygen were varied from 21 to 1140 kPa, as summarized in Table 11. Preliminary experimenta carried out at 100 OC in the stirred reactor showed that within a 24-h period the peroxide formation in fuels 2-4 was substantial. On this basis, it was concluded that rates of peroxide formation measured in the 100-120 "C temperature range would be appropriate for predicting rates of peroxide formation at ambient conditions if the mechanism could be assumed to remain unchanged throughout the temperature range from ambient to 120 "C.
Results and Discussion Existing methods for the evaluation of the storage stability of distillate fuels include "bottle storage" under an atmosphere of air at 43 and 65 O C I O J 1 for extended periods ranging from weeks to months. To establish baseline data,
fuel 1 2 3
4 5 6
Table 11. Test Matrix oxygen pressure, kPa Ta'= 43* T = 65* T = 80 T = 100 T = 120 "C "C OC OC "C 21 240 790, 1140 240,790 240 790, 1140 240,790 21 21 790, 1140 240,790 240 240 790, 1140 240,790 21 21 790,1140 240,790 240 240 21 790,1140 240,790
"Stress temperature. bBottle storage testa using ambient air; all other testa under oxygen pressure between 240 and 1140 kPa abs.
bottle storage tests were performed on all six model fuels. As shown in Table 11, fuels 1-4 were aged at 43 "C, and fuels 6 and 6 were aged at 65 "C. The results of these bottle storage tests were compared with more rapid hightemperature fuel oxidation tests at 80, 100, and 120 "C performed in the stirred pressure reactor. The rate of peroxide formation varied with fuel type and temperature, but the characteristics of the peroxide formation process were similar in all of the test fuels. Peroxides formed slowly in the early stages of oxidation. Later the rate became considerably higher as the peroxide concentration in the fuel increased. It became apparent that the rate of peroxide formation depended on the concentration of peroxides in the fuel. As the peroxide concentration increased, the rates of oxygen consumption and the formation of gums and water also increased. Analysis of the fuel for gums and water content was considerably less accurate than that for peroxide content. However, the formations of gums and water seemed to increase in direct proportion with the concentration of peroxides in the fuel. These products were particularly evident in the more reactive fuels, such as 4, which exhibited much higher rates of peroxide formation and thus produced more easily measured amounts of gums and water. It is well-known that the autoxidation of hydrocarbons is based on a free-radical mechanism,12which includes the familiar radical initiation, propagation, and termination reaction steps. While several reaction steps are conceivable in the overall autoxidation of hydrocarbon fuels, the formation of alkyl hydroperoxides, ROOH, water and gums may be described by the mechanism shown in reactions 1-7. RH + 0 2 ROOH ROOH
- + HO'
RO'
+ RH R' + HzO RO' + RH -,R' + ROH R' + 0 2 ROz' ROz' + RH R' + ROOH R02' + R02' products HO'
4
(3) (4)
-+
In this mechanism, the alkyl peroxide, ROOH, itself initiates the chain mechanism defined by reactions 2-7. Reaction 1 represents a relatively slow global process that may possibly consist of several unknown elementary processes. It has been included in the mechanism to simply account for the formation of the trace amount of peroxide required to initiate reactions 2-7. Free radicals, HO' and ROO, are formed by the decomposition of ROOH in reaction 2. These radicals react (12)Frost, A. A.; Pearson, R. G . Kinetics and Mechanism; Wiley: New York, 1961; pp 248-251.
Peroxide Formation in Jet Fuels rapidly, forming alkyl radicals, R', in reactions 3 and 4. Reaction 3 probably accounts for a significant amount of the water formed in the oxidation. In the presence of oxygen, the alkyl radicals, R', are rapidly converted to alkylperoxy radicals, ROz', via reaction 5. RO; is a relatively stable free radical because reaction 6 is slow compared to reactions 3-5.12 Since R02' reacts slowly, it tends to build up, and its concentration is much higher than that of the other free radicals. Consequently,the radical pool is depleted principally by the recombination of ROz' in reaction 7. Mayo and L a d 3 have suggested that freeradical termination reactions are responsible for the formation of gums in the autoxidation of hydrocarbons. This suggestion seems to be in agreement with the experimental results that show the rate of gum formation increases in proportion with the buildup in peroxide concentration. In a pristine fuel completely devoid of ROOH, the chain reaction will not start until a trace of ROOH is formed by another process denoted by reaction 1. Since the global reaction 1is relatively slow and contributes very little to the bulk formation of ROOH, it may be neglected in the expression for the rate of formation of ROOH, as shown in eq I. d[ROOH]/dt k,[RO,'][RH] - k,[ROOH] (I)
If the steady-state approximation is made for the freeradical concentrations [HO'], [RO'], [R'], and [RO,'], eq I may be expressed as d[ROOH]/dt = k~(2k~/k7)'/2[RH][ROOH]'/2 - k,[ROOH] (11)
Energy & Fuels, Vol. 2, No. 6, 1988 731 40
30
Ba fi 20 n z
z
p"
10
v
0 0
,
[ROOHI'/' k5(k2/2k,)1/2[02]t (V) where [O,] is the concentration of oxygen dissolved in the fuel. The results of the bottle storage experiments at 43 and 65 OC are shown in Figures 1and 2. There is favorable correlation between the square root of the peroxide concentration and the stress duration. If it is assumed that the rate of peroxide formation is independent of the oxygen concentration, the slopes of the lines in Figures 1 and 2 may be expressed as global rate constants, k, given by k N k6(k2/2k7)1/2[RH] where [RH] is assumed to be constant. The actual concentrations of reactive RH molecules in a fuel are not (13)Mayo, F.R.;Lan, B.Y . R e p r . Pap.-Am. Chem. SOC.,Diu. Fuel Chem. 1983,28(6), 1209-1216.
2000
3000
4000
STRESS DURATION, HOURS
Figure 1. Oxidation of fuels 1-4 under air at 43 O C . 30-
s
n
20
-
p.
v i 2
W
g
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W
a U n -
0
Integration of,eq I1 over the limits of [ROOH] = 0 to [ROOH], and t = 0 to t gives in (1- kz/k6(k7/2k2)'/2[ROOH]1/2/[RH]) = -k2t/2 (111) which may be written in approximate form as [ROOH]1/2"P kg(k~/2k7)'/~[RH]t (IV) Equation IV shows that the formation of peroxides is independent of the oxygen concentration and appears to depend only on the hydrocarbon concentration, which can be assumed to be essentially constant in an autoxidation process. It is important to notes that, if the partial pressure of oxygen is too low Oess than about 7 kPa), reaction 5 may become a rate-controlling step in the mechanism, and reaction 7 would then be replaced by reaction 8. R'+R'-Rz (8) For the oxygen-starved reaction, it can be shown that the peroxide concentration may be expressed as
1000
200
400
800
800
1000
STRESS DURATION, HOURS
Figure 2. Oxidation of fuels 5 and 6 under air at 65
OC.
known, but the fact that the curves in Figures 1 and 2 remain linear even after substantial amounts of peroxides are formed supports the assumption that only a negligible portion of the total RH is consumed in the autoxidation process. The results in Figure 1 show that the susceptibility of the fuels to form peroxides a t 43 O C increases in the following order: fuel 1 < fuel 3 < fuel 2 < fuel 4. Clearly, fuels 1 and 3 are quite stable compared to fuels 2 and 4. Aside from being the most reactive, fuel 4 waa different from the others in that it exhibited an induction period that lasted about 4 weeks. Since fuels 3 and 4 are the same except for the fact that fuel 4 was alumina treated to remove possible polar compounds, it seemed that fuel 3 contained some form of oxidation inhibitor. The induction period observed in fuel 4 appears to be caused by the remnants of the inhibitor that may have been present in fuel 3 before alumina treatment. This invites the thought that the results shown in Figure 1for fuel 3 may represent an induction period. Although a trace analysis of fuel 3 failed to show known antioxidants such as phenols and amines, this fuel most probably contained some form of radical scavenger. Since alumina treatment removes only the most polar compounds from the fuel, present in minute amounts, it is unlikely that it would have an influence on the bulk properties. Relatively low concentrations of inhibitors tend to reduce the rate of oxidation by lowering the steady-state radical concentration, which basically is the same as increasing the radical termination reaction rate. The bottle storage experiments on fuels 5 and 6 were carried out at 65 OC. Figure 2 shows the average results of three independent measurements with an induction period lasting about 4 weeks, followed by a more rapid buildup of the peroxide concentration in fuels 5 and 6. The
Fodor et al.
732 Energy & Fuels, Vol. 2, No. 6, 1988
ICNa31 -0- FUEL 1
-a-
FUEL 2
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4
j
.
fi 20 -
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50
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STRESS DURATION, HOURS
Figure 3. Oxidation of fuels 1-4 under 790 and 1140 kPa of oxygen at 80 "C.
I
0
5
10
15
5
10
15
STRESS DURATION, HOURS
20
25
STRESS DURATION, HOURS
Figure 4. Oxidation of fuels 1-4 under 240 and 790 kPa of oxygen at 100 O C .
induction periods give strong correlationsof the square root of the peroxide concentration with test duration. The slopes of the post-induction periods are not well resolved because the end of the induction period is not defined, but, nevertheless, they represent minimum values of the global rate constants. It is possible only to speculate on reasons why fuels 4-6 show induction periods. The most likely explanation seems to be that fuels 4-6 are similar to fuel 3, which shows only an induction period. Compared to fuel 3, fuels 4-6 also appear to contain small amounts of a radical scavenger that, in effect, increases the radical termination reaction rate, thus reducing the free-radical concentration and lowering the rate of oxidation. Alumina treatment of fuels 5 and 6 was not undertaken, and the effect of radical scavengers on peroxide inhibition was not investigated in this study. However, previous studies13on the autoxidation of hydrocarbons have shown that similar induction periods appear in autoxidations of hydrocarbons when antioxidants are added. The durations of these induction periods were found to be proportional to the amount of antioxidant added to the hydrocarbon. Figures 3-5 show the results of peroxide formation in the stirred pressure reactor for fuels 1-4 at 80, 100,and 120 "C, respectively. In these experiments, the partial pressure of oxygen in the reador was varied to determine if the rate of peroxide formation was dependent on the amount of oxygen dissolved in the fuel. The measurements were made by using oxygen partial pressures of 790 and 1140 kPa at 80 "C, and oxygen partial pressures of 240 and 790 kPa at 100 "C. In both of these experiments, there were approximately an equal number of data obtained at the two oxygen concentrations. I t is apparent in Figures 3 and 4 that the variation in the partial pressure of oxygen
Figure 5. Oxidation of fuels 1 4 under 240 kPa of oxygen at 120
"C.
had no effect on the rate of peroxide formation. This agrees with other studies? which found that the rate of autoxidation does not become oxygen limited until the partial pressure falls below 7 kPa. It is important to note, however, that the lower limit of 7 kPa is an approximate value, which no doubt depends on the temperature and the reactivity of the fuel with oxygen. Substantially greater concentrations of oxygen may well be required in autoxidation reactions carried out at higher temperatures and on fuels that are highly reactive. Recent results of Watkins, et d.,'on peroxide formation in jet fuels at 43-100 "C seem to be indicative of an oxygen-limited autoxidation process. Their experiments were carried out in sealed brown-glass 500-mL bottles filled with 300 mL of fuel. The oxygen content of the vapor space was not analyzed during the course of the autoxidation to determine if the reaction rate might be limited by the oxygen concentration in the fuel. Nor was the fuel stirred to increase the rate of oxygen dissolution, which is otherwise a relatively slow diffusion process. The results of this work showed peroxides building up quite rapidly at first and then tapering off and actually decreasing as time progressed. In the higher temperature experiments, the buildup in peroxide concentration began to decrease much earlier in the reaction and at a lower peroxide concentration. Basically, the peroxide concentration is a balance between the rate of formation and the rate of decomposition. When the reaction is oxygen-concentration limited, the decomposition process will eventually dominate. The effect of temperature on peroxide formation in jet fuels is quite apparent when the low-temperature bottle storage results are compared with the higher temperature stirred pressure reactor data. It takes several months at 43 "C to form the same amount of peroxide that is formed in less than 10 h at 120 "C. In the 80-120 "C temperature range, the susceptibility of fuels 1-4 to form peroxides increases in the following order: fuel 1 < fuel 2 < fuel 3 < fuel 4. Contrary to the measurements made at 43 "C, it was found at higher temperatures that fuel 3 was more reactive than fuel 2. It will be shown later that this change in the order of reactivity is due to a higher activation energy for the formation of peroxides in fuel 3. Figures 6 and 7 show the formation of peroxides from fuels 5 and 6 in the stirred pressure reactor at 100 and 120 "C,respectively. These data also show that the rate of peroxide formation is significantlyincreased over the bottle storage results obtained at 65 "C. While the induction period lasted for about 4 weeks at 65 "C, the corresponding times at 100 and 120 "C were reduced to 23 and 2 h, respectively. The fact that the induction period decreases as the temperature is raised, suggests that it is caused by
Peroxide Formation in Jet Fuels
Energy & Fuels, Vol. 2, No. 6,1988 733 4 -
+/+
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Figure 6. Oxidation of fuels 6 and 6 under 240 kPa of oxygen at 100 O C .
Figure 8. Arrhenius plot of the oxidation of fuels 1-4. 2 -El-
-2
-
2.5 0
2
4
6
e
10
STRESS DURATION, HOURS
Figure 7. Oxidation of fuels 6 and 6 under 240 kPa of oxygen at 120 O C . Table 111. Linear Regression Analysis of the Arrhenius Plots Based on In k = In A - E J R T fuel 1 2 3 4 5 6 6 6 a
temp E,, kcal/ Deriod ranee. O C mol 43-120 19.4 unknowna 43-120 21.6 unknowna 29.4 43-120 unknowna post-inductionperiod 43-120 22.0 65-120 23.0 induction period 21.3 post-inductionperiod 65-120 65-120 21.9 induction period 21.5 post-inductionperiod 65-120
1nA 23.5 29.4 40.3 30.5 30.5 28.8 28.2 28.8
R2 0.963 0.998 0.985 0.998 0.999 0.999 0.969 0.997
Within experimental time limits.
a radical scavenger in the fuel. As the reaction rate increases, the steady-state concentration of free radicals must also increase. With higher free-radical concentrations in the fuel, the free-radical scavengers become depleted much more rapidly so the induction period is shortened. The mechanism proposed in reactions 2-7 accounts for the formation of peroxides throughout the temperature range 43-120 O C because the data correlate favorably according to eq IV and the global rate constants correlate with the reciprocal of the absolute temperature. Figures 8 and 9 show Arrhenius plots of the global reaction rate constants obtained for fuels 1-4 and fuels 5 and 6, respectively, and Table I11 gives the correlation parameters, which are based on least-squares fits of the data. The linearity of the Arrhenius correlations support the conclusion that the reaction mechanism remains unchanged over the temperature range of 43-120 "C. It is also interesting to note that the global rate constants obtained from the bottle storage experiments at 43 and
FUEL 5 POST IND. PERIOD
2.8
2.7
2.0
2.9
3
(~/T)X~O~
Figure 9. Arrhenius plot of the oxidation of fuels 6 and 6.
65 "C correlate favorably in the Arrhenius plots with the rates measured at higher temperatures in the pressure reactor. Since the partial pressure of oxygen was about an order of magnitude less in the bottle storage experiments than it was in the stirred pressure reactor experiments, the results of the Arrhenius plots give further confirmation of the oxygen independence of the autoxidation reaction and show that the fuels were supplied with adequate oxygen in the bottle storage experiments. The activation energies given in Table I11 for peroxide formation in the test fuels tend to fall into the 19-22-kcal range except for fuel 3. The activation energy of fuel 3 is about 7 kcal/mol greater than that of fuel 4. Since the only difference between fuels 3 and 4 is alumina treatment, it is surprising that there is such a marked change in the activation energy. It does not appear that the inhibiting substance removed from fuel 3 by alumina treatment was an antioxidant such as a substituted phenol. Antioxidants are not expected to increase the activation energy because they tend only to influence the radical termination rate. I t is concluded in the present study of jet fuel autoxidation that the tendency to form peroxides depends on both the reactivity of the bulk hydrocarbons and the presence of small, but significant, concentrations of natural inhibitors. If it is argued that the rate-controlling step in the autoxidation process is the attack on the C-H bond by the R02' radical, i.e., reaction 6 in the proposed mechanism, then the ease of hydrogen abstraction determines the overall rate of ~xidation.~J"" Since the order (14)Betta, J. Q. Rev. 1971,25, 265-288. (15)Szwarc, M.;Buckley, R. P.; Greaser, J. A.epr.-Am. Chem. Soc., Diu.Pet. Chem. 1958,3(1), 149-159. (16)Korcek, S.;Chenier, J. H. B.; Howard,J. A.; Ingold, K. U. Can. J . Chem. 1972,50, 2285.
734
Energy & Fuels 1988,2,734-739
of free-radical stability and attack on the carbon-hydrogen bond is benzylic and allylic H > tertiary H > secondary H > primary H, it was expected that fuels with different oxidative tendencies would exhibit gross differences in their NMR spectra. To examine this possibility, the IH and 13Cnuclear magnetic resonance (NMR) spectra of fuels 1,2, and 4 were measured. The proton-type assignments were made according to the recommendations of Netzel and Hunter.18 Surprisingly, the relatively large difference in the reactivity with oxygen was not reflected in the NMR spectra of fuels 1,2, and 4. The spectra did not show any significant enhancement of the more reactive C-H bond types in fuel 4 compared to those in fuels 1 and 2. Apparently these differences in the strengths of C-H bonds in the test fuels are more subtle than expected and cannot be resolved by NMR spectroscopy. The absence of correlation is probably due in part to the complexity of fuel composition. However, it is also possible that the concentration of the active species (RH) in these fuels is relatively low. It was argued in the proposed mechanism that RH is in high concentration, based on the relative amount of peroxide formed. When it is realized that measured peroxide concentrations seldom exceed 3000 ppm, then an RH concentration of as little as 2-3% of the fuel could be considered to be relatively high. Yet, a species in the 2-3% concentration range is at the threshold of detection by NMR. These questions can only be answered by continued experimental work on the kinetics of autoxidation in jet fuels.
Conclusion The rates of peroxide formation in six model kerosenes were measured in the temperature range 43-120 OC with oxygen partial pressures ranging from 10 to 1140 kPa. To explain the rate of peroxide buildup in the fuels, a kinetic model of the autoxidation process was developed on the (17) Nixon, A. C. In Autoxidation and Antioxidants; Lundberg, W . O., Ed.; Interscience: New York, London, 1962; Chapter 17. (18) Netzel, D. A.; Hunter, P. M. DOE/LETC/RI-81-1, Laramie, WY,
May 1981.
premise that peroxide decomposition is the principal free-radical initiation step. In accordance with this model, it was found that the square root of the peroxide concentration was proportional to the stress duration. Global rate constants determined from the peroxide concentrationtime histories were independent of the partial pressure of oxygen but strongly dependent on the stress temperature. Arrhenius correlations of the global rate constants showed that the mechanism of peroxide formation remained unchanged in the temperature range 43-120 OC. The activation energies of the fuels ranged from 19 to 22 kcal/mol except that for fuel 3, which was 29 kcal/mol. Because there was a significant variation in the activation energies of peroxide formation, it is concluded that peroxide potential can not be predicted from single-point rate measurements at elevated temperatures. However, the results of this work encourage the development of a timely test method that predicts rates of peroxide formation at ambient conditions when two or more measurements are made at elevated temperatures.
Acknowledgment. This work was conducted at the Belvoir Fuels and Lubricants Research Facility, located at Southwest Research Institute (SwRI), for the Naval Research Laboratory (NRL) under Contract No. N00014-85-C-2520. For completeness, this paper includes the results of earlier work funded by the Naval Air Propulsion Center (NAPC). NAPC funding was provided by a Military Interdepartmental Purchase Requisition through the US.Army Belvoir Research, Development and Engineering Center under Contract DAAK70-85-C0007, with F. W. Schaekel as Contracting Officer's Representative. The project monitors for the Navy were Dr. D. R. Hardy of NRL and C. J. Nowack, G. E. Speck, and Lynda C. Turner of NAPC. We gratefully acknowledge the invaluable discussions and participation by colleagues G. H. Lee, 11, W. D. Weatherford, Jr., N. F. Swynnerton, Marilyn Voigt, K. B. Jones, J. J. Dozier, and Deborah Toles. Editorial assistance provided by J. W. Pryor, Marilyn Smith, Sherry Douvry, and LuAnn Pierce is also acknowledged with appreciation.
Analysis and Prediction of Product Distributions of the Fischer-Tropsc h Synthesis Timothy J. Donnelly, Ian C. Yates, and Charles N. Satterfield* Department of Chemical Engineering, Massachusetts Institute of Technology, Cambridge, Massachusetts 02139 Received March 14, 1988. Revised Manuscript Received June 18, 1988
A method is developed for calculating the three parameters needed to characterize the carbon number distribution of products of the Fischer-Tropsch synthesis. Experimental data are fit by a modified Schulz-Flory model that has two chain growth probabilities, using nonlinear regression. Excellent fit is shown for data from precipitated iron and fueed magnetite catalysts. The model is used to calculate selectivity information of interest in catalyst comparison and reactor design. Advantages of this method over asymptotic regression methods are discussed in detail. Introduction The r>roducts of the Fischer-Tror>sch synthesis are primariiy linear hydrocarbons distributed -over a wide range of carbon numbers. Herington' reported that a
model of stepwise addition of single-carbon units could predict the fraction of product at each carbon number. (1) Herington, E. F. G. Chem. Ind. (London) 1946, 65, 346.
0887-0624/88/2502-0734$01.50/00 1988 American Chemical Society
