Ind. Eng. Chem. Res. 1995,34, 2285-2291
2285
Oxidation of n-Butane: Transition in the Mechanism across the Region of Negative Temperature Coefficient Richard D. Wilk* Department of Mechanical Engineering, Union College, Schenectady, New York 12308-2311
Richard S. Cohen Department of Mechanical Engineering, Temple University, Philadelphia, Pennsylvania 19122
Nicholas P. Cernansky Department of Mechanical Engineering, Drexel University, Philadelphia, Pennsylvania 19104
An experimental investigation has been conducted of the transition i n the oxidation chemistry of n-butane across the region of negative temperature coefficient from low to intermediate temperatures. The experimental results indicated a region of negative temperature coefficient between approximately 640 and 695 K and a shifi in the nature of the reaction intermediates and products across this region. On the basis of these experimental results and some recent kinetic modeling results for the reaction R 0 2 * ROz, a new mechanism is presented for n-butane to describe the observed phenomena. At low temperatures the major reaction path of
+
butylperoxy is isomerization followed by 0 2 addition, further isomerization, and decomposition to mainly carbonyls and OH radicals. At intermediate temperatures, the major reaction of butylperoxy is isomerization followed by decomposition to butenes and HOz and to epoxides and OH. The mechanism is consistent with these and other hxperimental results and predicts the negative temperature coefficient and the shift in product distribution with temperature.
Introduction Better descriptions of the complex processes of preflame hydrocarbon oxidation chemistry are necessary for adequate modeling of many practical and laboratory combustion devices. For example, several experimental and modeling studies on autoignition in relation to engine knock (Cernansky et al. 1986; Leppard, 1987; Green et aZ.,1987; Henig et al., 1989) have focused attention on the importance of preignition chemistry and the resulting heat release to the autoignition process. Also, it has been demonstrated that the propensity of a fuel to autoignite is related to its tendency t o form cool flames. Consequently, there is an incentive to develop a fundamental understanding of the oxidation mechanism of hydrocarbon fuels at temperatures which are characteristic of those preceding autoignition, i.e., below about 800-900 K. Previously, all of the chemistry below 900 K was classified as low temperature or cool flame chemistry. This generalization proved to be inaccurate since, actually, there are at least two distinct reaction regimes for aliphatic hydrocarbons below 900 K a low temperature regime below about 650 K and what we have designated as an intermediate temperature regime (Wilk et al., 1986) which extends from about 650 K up to about 900 K. The oxidation chemistry in each of these reaction regimes is quite different. In addition, the chemistry in each of these regimes is very different from that occurring above about 900 K. The differences in the chemistry can be classified by the different dominant radical species, the stable intermediate and product species formed, and the overall reaction rate. In addition, the different mechanistic reaction paths which occur in the different temperature regimes and especially the transition in the mechanism between tem-
* Address correspondence to this author.
perature regimes are responsible for many very interesting combustion phenomena. These two temperature regimes are usually separated by a region of negative temperature coefficient (NTC). This is a phenomenon characteristic of aliphatic hydrocarbon oxidation in which the overall reaction rate decreases with increasing temperature. This reduction in the rate occurs over a temperature range on the order of 100 deg (usually occurring between 600 and 700 K). This region represents a transition in the oxidation chemistry from the low temperature to intermediate temperature reaction regime, where the rate of the accelerating or branching mechanism of the low temperature regime begins to subside before that of the intermediate temperature regime becomes important; thus the overall reaction rate decreases. It should be noted that the temperature boundaries separating these reaction regimes are dependent on the reactant concentrations and the total pressure. All of the temperature boundaries mentioned above are for low pressures (on the order of 1 atm). Increasing the pressure will likely shift the upper temperature limits of the low and intermediate temperature regimes to higher temperatures (Dryer, 1991). A detailed understanding of the oxidation chemistry in each of these regimes as well as the transition from one regime to the other is important, as noted, because it will allow us to adequately account for preflame processes and to determine their contribution to autoignition and their significance in practical systems. n-Butane has already been the subject of numerous experimental and modeling studies. Detailed reaction mechanisms have been developed for n-butane by Cathonnet et ul. (19811, Warnatz (19831, Pitz et ul. (1984, 1986), and Chakir et al. (1989). These models have been shown t o be accurate in the high temperature regime by reproducing results from high temperature experiments such as flow reactors, shock tubes, and jet stirred
0888-5885/95/2634-2285$09.00/00 1995 American Chemical Society
2286 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995
reactors. The model of Pitz et al. (1984) was modified and extended to incorporate reactions which are important at low and intermediate temperatures and utilized subsequently to model autoignition engines, rapid compression machines, and two-stage flames (Cernansky et al., 1986; Green et al., 1987; Pitz et al., 1988; Carlier et al., 1990; Corre et al. (1992); Minetti et al., 1994). A new mechanism was developed recently by Kojima (1994) and used to model autoignition delay times from a shock tube and from a yapid compression machine. A major conclusion resulting from many of these studies was that more work is required to elucidate the details of the mechanism occurring at low and intermediate temperatures, especially the process by which the mechanism shifts between these two regimes. A great deal of experimental work over the years has been devoted t o the study of the autoignition and preflame processes of n-butane (Bardwell, 1955; NanChiang and Bardwell, 1955; Cherneskey and Burdwell, 1960;Salooja, 1967; Euker and Leinroth, 1970; Dechaux et al.,1971; Carlier and Sochet, 1975; Leppard, 1987; Sahetchian et al.,1990; Griffiths and Nimmo, 1985; Griffiths et al., 1993; Chandratna and Griffiths, 1994). Also, there have been a number of studies recently which examined the autoignition of n-butane in blends with other compounds, such as isobutane, hydrogen, methane, peroxide, and NO (Henig et al., 1989; Refael and Sher, 1990; Griffiths et al., 1990; Inomata et al., 1990; Bromly et al., 1992). Fundamental studies examining some of the details of the oxidation mechanism of n-butane have also been made (Baker et al., 1975a,b; Berry and Cullis, 1970; Berry et al., 1970; Sheinson and Williams, 1973, 1984). From these studies additional reaction paths, believed to be important in the low temperature n-butane mechanism, in addition t o some important reaction rate parameters, were determined. Among these was the intramolecular isomerization and decomposition path of the butylperoxy radicals. Despite the enormous body of work that has been done to elucidate the n-butane mechanism, there remain some areas which are still unclear. One of these has t o do with the transition in the mechanism from low to intermediate temperatures and the associated occurrence of the negative temperature coefficient. The work described in this paper complements these many previous studies on n-butane by identifying the region of negative temperature coefficient and focusing on the underlying chemical mechanism responsible for the transition in the chemistry across this region from the low temperature/cool flame regime to the post NTCI intermediate temperature regime. A new explanation for the n-butane system is offered based on the experimental results and some recent new thinking regarding the alkene-forming reactions of alkyl and 0 2 .
Experiment The experimental study of n-butane oxidation was carried out using a conventional static reactor system. Details of the experimental facility have been described previously (Wilk et al., 1986,1987);thus only the main features will be briefly reviewed here. The system consists of a cylindrical Pyrex reaction vessel (volume 1395 cm3, diameter 10 cm, surface-tovolume ratio (SA9 0.5 cm-l) located inside a temperature-controlled compartment. The pressure inside the vessel was measured with a Setra Model 204 pressure transducer. The temperature a t the center of the vessel was measured by a pt/pt-13% Rh thermocouple con-
620
I K
:=E74
" I "
0
100
200
300
400
500
600
Time (s)
Figure 1. Pressure-time profiles for n-butane oxidation. P c ~ H ~ ~ = 23.6 Torr; Po, = 47.1 Torr; P N = ~ 479.3 TOIT.
structed of 0.05 mm diameter wires and coated with a thin layer of silica to prevent localized catalytic heating. The reactants were first premixed in a separate vessel and then rapidly admitted to the preheated, evacuated reaction vessel. At a selected time during the course of the reaction a portion of the reactor contents was withdrawn into an evacuated sampling loop quenching the reaction. Stable species analyses were conducted using a Varian 3700 gas chromatograph (GC) equipped with a flame ionization detector (FID). A gas sampling valve was used to introduce a portion (0.25 m3)of the acquired sample into the GC. The sample was separated on a 2.4 m x 3mm, 80/100 mesh Porapak Q column using helium carrier gas (20 cm3/min). A post column nickel catalyst was used to methanize CO and COz and facilitate their detection with the FID. Temperature programming (4 min at 36 "C; 10 C/min to 180 "C;20 min at 180 "C)was used to obtain adequate separation of the stable hydrocarbons and carbon oxides. For this work, oxygen (99.8% pure), nitrogen (99.999% pure), and n-butane (99.5% pure) were used as the reactants. Experiments were carried out for a mixture of n-C1HldOflz: 1/2/20.333molar ratio (4 = 3.25), at an initial total pressure of 550 Torr. The initial reaction temperature was varied from 554 to 737 K encompassing the NTC region and portions of each of the low and intermediate temperature regimes.
Results and Discussion Phenomenological Behavior of the Overall Rate of Reaction. Selected measured pressure histories for the oxidation of n-butane are presented in Figure 1.As shown, the reaction times for the process at these temperatures are relatively long, on the order of tens to hundreds of seconds. The pressure profiles exhibit the basic S-shape characteristic of autocatalytic reactions. The profiles for initial temperatures of 574 and 563 K display sharp pressure pulses on top of the S shape. These indicate the occurrence of cool flames. The phenomenology of cool flames in low temperature hydrocarbon oxidation has been the subject of many studies and is reviewed extensively in the literature (Sokolik, 1963; Fish, 1968a; Lignola and Reverchon, 1987). At the conditions of these experiments with
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2287 250
200 CI
c
'ee
-B
150
c
x
Q
100
E
n
c
s
s
50
0 900
550
600
650
700
750
;OO
550
600
650
700
750
Initial Temperature (K) Figure 2. Effect of initial temperature on the induction period H 23.6 ~ ~ Torr; PO,= 47.1 Torr; P N = ~ for n-butane oxidation. P c ~ = 479.3 Torr.
Figure 3. Effect of initial temperature on the maximum rate of pressure rise for n-butane oxidation. P c ~ = H23.6 ~ ~ Torr; Po,= 47.1 Torr; P N=~ 479.3Torr.
n-butane, cool flames were observed at temperatures ranging from 563 t o 633 K. The number of cool flames increased from one at 563 K to a maximum of three at 596 K and then decreased back to one at 633 K. The maximum changes in temperature and pressure attained during a cool flame pulse reached about 32 deg and 60 Torr, respectively. Historically pressure-time profiles have been used to infer some information about the kinetics. In constantvolume systems the pressure profile often mirrors the fuel disappearance profile. Thus the rate of change of pressure is indicative of the overall reaction rate. From the pressure profiles in Figure 1at initial temperatures of 664 and 715 K, it can be seen that the reaction rate increases slowly at first, accelerates to a maximum, and then decreases again resulting in an overall pressure increase for the reaction, AP. Two parameters can be obtained from the pressure-time profiles which are often used as indicators of the reactivity or ignitability of a mixture. The induction period, z, is defined as the time period from the admission of the fuelloxidizer mixture to the onset of rapid pressure rise. It is analogous to the ignition delay time in systems where autoignition occurs. The kinetic interpretation for induction period is the initial time required for buildup of a sufficient pool of radicals in the system which can then rapidly accelerate the oxidative process. The other parameter used is the maximum rate of pressure rise, (dP/dt),,, which is essentially the inflection point in the pressure profile. It can be seen from Figure 1that the overall rate of reaction the case To = 664 K is considerably slower than that a t To= 574 K and To= 563 K and has a longer induction period. This decrease in the overall rate with increasing temperature is indicative of a negative temperature coefficient (NTC). Thus the condition at 664 K lies within the region of negative temperature coefficient. The region of negative temperature coefficient for n-butane was determined by plotting the induction period over the range of initial temperatures. The effect of initial reaction temperature on induction period is shown in Figure 2. The induction period decreases exponentially up to approximately 640 K and at 655 K begins a dramatic increase of 2 orders of magnitude over a 40 deg range from 655 to 695 K. It should be noted
that the determination of the beginning of the NTC region may be subject to some uncertainty. Between 620 and 655 K the induction periods are very short (on the order of seconds) and there is the possibility that the mixture is beginning to react as it is still in the process of entering the reactor. Nevertheless, the measured induction periods actually begin to increase at approximately 640 K. Finally, above 695 K, the induction period again decreases with increasing initial temperature. Thus, based on the induction period data, for the initial pressure and concentration conditions of these experiments, the region of NTC was established between 640 and 695 K. This result is generally consistent with the results obtained by Nan-Chiang and Bardwell (1955), Salooja (19671, and Dechaux et al. (1971). We also examined the effect of initial temperature on the maximum rate of pressure rise. This resulted in a plot (Figure 3) that was nearly the inverse of the induction period plot. The maximum rate increased, decreased, and then increased again with increasing initial temperature. The temperature span of the NTC region was about the same as that obtained from the induction period data (55-60 deg). However, the major difference was that the appearance of the region was shifted to lower temperatures by approximately 25 deg. This offset could be due, in part, t o the uncertainty in the induction period measurements between 620 and 655 K. It may also indicate that there may be some fundamental differences in the ways in which induction period and maximum rate of pressure rise are representative of the overall reactivity. Outside of the NTC ranges indicated by the induction period and maximum rate of pressure rise data, there is a high degree of correlation between the reciprocal induction period and maximum rate of pressure rise in each of the low and intermediate temperature regimes. Figures 4 and 5 show plots of l/z against (dP/dt)m, for the low and intermediate temperature regimes, respectively. Only the data that were truly shown to be in the low and intermediate temperature regimes for each case were included in the plots. The maximum rate of pressure rise, (dP/dt)m,, can be said to be related to an overall reaction rate coef-
Initial Temperature (K)
2288 Ind. Eng. Chem. Res., Vol. 34,No. 7, 1995 Table 1. Overall Activation Energies for n-Butane Oxidation
0.20
E1 (kcaL"o1)
low temperature intermediate temperature
0.15
Ez (kcal/mol) 41.0 28.8
induction period with initial temperature is often expressed by a form similar to that of eq 1:
A
k
39.4 33.4
0.10
P r
\
0.00
0
20
40
60
80
100
120
140
(dP/dt)max (torr/min)
Figure 4. Correlation of reciprocal induction period and maximum rate of pressure rise for n-butane oxidation (low temperature regime). P c ~ = H 23.6 ~ ~ Torr; Po, = 47.1 Torr; PN,= 479.3 Torr. 0.020
0.015
c b, Y
P
indicates how sensitive the induction period is t o initial temperature and is determined from the slope of a fitted line through the data plotted as In z against reciprocal temperature. The data were applied to eqs 1 and 2, and overall activation energies were determined for each of the temperature regimes. These are presented in Table 1. The values of E1 and Ea for the low temperature regime were fairly close to each other. The E1 and E2 values for the intermediate temperature regime were not as close but both were less than their respective low temperature values. Thus, the overall rate of reaction is more sensitive to temperature in the low temperature regime. Chemical Species Measurements. In order t o assess the changes which occur in the chemistry across the NTC region for n-butane, species measurements were made a t temperatures in both the low and intermediate temperature reaction regimes and compared. These data are shown in Figures 6 and 7 as species mole fraction plotted against time. Also shown are the pressure and temperature histories for the reactions. Actual data points are plotted along with smoothed curves. The oxidation of n-butane yielded several reaction products. The primary products included carbon monoxide, carbon dioxide, butene, propene, ethene, methanol, methane, formaldehyde, acetaldehyde, acrolein, and propionaldehyde. Minor products included ethane, ethanol, acetone, butene oxides (2,3epoxybutane (2,3-C4H80) and 1,2-epoxybutane (1,2C4H&)), butyraldehyde, 2-butanone, pentane, methacrolein, tetrahydrofuran, and crotonaldehyde. Although most of the species were able to be quantified using the GC, column limitations made it such that some of the species had t o be combined as single quantities. Chief ,among these were the butenes. 1-Butene, cis-2-butene, and trans-2-butene could not be individually separated, so they had to be represented as a single combined quantity. Other low concentration species also had to be represented as single combined quantities: acrolein and propionaldehyde; 1,2-epoxybutane, butyraldehyde, and 2-butanone. The high concentrations of carbon monoxide relative to carbon dioxide should be noted. This is not surprising considering that the mixture is fuel rich. However, the richness of the fueVoxidizer mixture alone cannot account for the extremely large amounts of CO produced. It is due primarily t o the fact that we do not have complete combustion at these temperatures. The temperatures are simply too low to allow any significant conversion of CO to C02 to occur via CO + OH C02 H. This effect is also seen in the relatively small rise in temperature accompanying the reactions. The oxidation of CO to COz by the above step is one of the main exothermic processes in combustion. This step is not occurring t o any great extent a t these conditions, and hence the low exothermicity observed.
E2
0.05
0.010
\ r
0.005
O.Oo0
0
10
20
30
40
(dP/dt)max (torrhin)
Figure 5. Correlation of reciprocal induction period and maximum rate of pressure rise for n-butane oxidation (intermediate temperature regime). P c ~ = H 23.6 ~ ~ Torr; PO,= 47.1 Torr; PN,= 479.3 Torr.
ficient, kov. Therefore, the variation in (dP1dt)mS with initial temperature over a specific temperature range can be expressed by an Arrhenius-type equation of the form
where C1 is a preexponential constant and E1 represents an overall activation energy of the preflame reactions. The equation can be linearized and a regression analysis performed on the data. The overall activation energy is obtained from the slope of the fitted line through the data plotted as ln(dPldt),, against reciprocal initial temperature. E1 is an indicator of how sensitive (dPl dt)" is t o temperature over a specified temperature range. The induction period is representative of a characteristic reaction time related to llkov. The variation of
+
-
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2289
"."
0.05
a
TO1.l
I"
r
.b
c/4
CH3OH
0.04
0.03
CH3CHO
t
0.02
0.005 0.01
0
200
100
300
O.OO0
100
0
200
300
200
300
Time ( 8 ) I
U.VUP
I
IC CH3COCH3
0.002
0.001
+ +
O.OO0
2.00.4 C3H7CHO CH3CHZCOCH3
2,3 C 4 H 8 0
0.ow 0
300
200
I00
Tlme
100
0
Time
(8)
(8)
CI
L.
0
c
Y
h3 M M
E
a
540
0
100
200
300
530
40U
Time (s) Figure 6. Species concentration profiles for n-butane oxidation. To= 563 K.P C , H = ~ ~23.6 Torr; Po2 = 47.1 Torr;
Comparison of the results in Figures 6 and 7 indicate a significant change in the important chemical pathways in going from the low to intermediate temperature regime. There is a general shift from mainly oxygenated hydrocarbon product species such as aldehydes and
= 479.3Torr.
alcohols at 563 K to primarily alkenes and alkanes at 715 K. At 563 K, formaldehyde and acetaldehyde are key products formed in nearly equal amounts, which is in agreement with the results of Leppard (1987). At 715 K, the maximum yield of each of these species
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 0.05
c2nY
0.04
C
c 0 .
L
-
0.010
0
e
0.03
U
e 4,
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400
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100
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400
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500
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(8)
620
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730
e
A
!5
-
580
- 705 - 700 - 695
570 560
550 540
0
710
100
200
300
400
690 500
Time (s) Figure 7. Species concentration profiles for n-butane oxidation. To= 715 K.P c , H = ~ ~23.6 Torr; PO,= 47.1 Torr; PN,= 479.3 Torr
decreased by over 50%. Methanol is the most abundant oxygenated hydrocarbon product formed (1.2%)a t 563 K. However, at 715 K, no methanol a t all was observed. On the other hand, the yields of methane, ethene, and propene formed at 563 K underwent a 10-fold increase
as the temperature was increased to 715 K. The total butene yield increased by a factor of 2 in going from 563 to 715 K. "his same general shift in the measured products with increasing temperature, from oxygenated to nonoxygenated products, was also observed in previ-
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2291 Table 2. Comparison of Reaction Intermediates and Products for n-Butane Oxidation at 50% Fuel Consumedo mol % species
co c02
CH4 C2H4 C2H6 C4Hs CH2O CH3CHO C3H6 CH30H C2H3CHO CzH5CHO CH3COCH3 2,3-C4HsO 1,2-C4HsO C3H7CHO CzH50H C5H12 tetrahydrofuran methacrolein crotonaldehyde
+
+
+ CzH5COCH3
563K 2.750 0.900 0.140 0.160 0.055 0.170 0.580 0.590 0.026 0.900 0.075 0.160 0.026 0.095 0.010 0.045 0 0 0
715K 2.150 0.720 0.350 0.650 0.005 0.400 0.250 0.240 0.310 0 0.130 0.058 0.020 0.079 0 0.013 0.026 0.010 0.002
aInitial Concentrations: P c ~ = H 23.6 ~ ~ Torr; Po2 = 47.1 Torr; P N=~ 479.3 Torr.
ous studies on propane and propene oxidation (Wilk et al., 1986, 1987, 1989). More of the initial fuel was consumed in the reaction at 715 K (78%)than at 563 K (58%). Therefore, in order to more accurately determine the effects of temperature on the product yields, it was necessary to compare the species data from each case on a more common basis. This was done by comparing the species concentrations at each temperature based a point where the fraction of n-C4Hlo consumed was the same. An example of these results is given in Table 2, which shows the product yields for each temperature taken at the point in the reaction at which 50%of the initial n-butane was consumed. On the basis of this comparison, the shift in the product distribution is still apparent. An exception to the trend of decreasing oxygenated hydrocarbons occurs with the unsaturated aldehydes acrolein, methacrolein, and crotonaldehyde. The yields of these species increase from 563 to 715 K. This behavior can be explained by the fact that these species are formed mainly from secondary oxidation reactions of the alkene intermediates (butene and propene). Thus, it can be expected that the yields of the unsaturated aldehydes will follow the behavior of the alkenes and increase with temperature. Primary Reaction Paths in the n-ButaneMechanism. In order to explain the observed differences in the product distributions, over the temperature range examined, and the occurrence of the NTC, it is necessary to consider the main reaction paths in the mechanism and determine how they are affected by the temperature. Traditionally, the transition in the chemistry and the product distribution accompanying the occurrence of the negative temperature coefficient have been attributed t o changes in the reaction between alkyl and 0 2 . The classical explanation for the transition was given by Benson (1965, 1982). Basically this involves addition of 0 2 to the alkyl radical (R) at low temperatures: R + 0,-R02
(3)
At low temperatures the RO2 abstracts H atoms from the fuel and aldehyde intermediates forming alkyl hydroperoxides. The alkyl hydroperoxides dissociate
into OH and alkoxy radicals (RO) which accelerate the reaction. At temperatures of 573-673 K, depending on the oxygen pressure, reaction -3 begins to compete with the forward association and the alternative path for the reaction of the alkyl radical and 0 2 is an H atom metathesis, leading to the production of the conjugate alkene and the hydroperoxy radical with increasing temperature:
R + 0,
-
alkene
+ HO,
(4)
Thus the OH in the system is replaced by less reactive HO2 resulting in a decrease in the reaction rate with increasing temperature, consistent with the observed NTC. The transition in the mechanism is then governed primarily by the equilibrium of the reaction R 0 2 RO2. Many kinetic studies though have discounted reaction 4 as being the correct mechanism for alkene formation (Slagle et al., 1984b; Bozzeli and Dean, 1990; Wagner et al., 1990; McAdam and Walker, 1987). Thus, the above parallel path, uncoupled mechanism for the reaction of R 0 2 has come under question recently and new thinking on the subject has emerged, motivated by the work of Slagle, Gutman, Wagner, and co-workers (Slagle et al., 1984a,b, 1986;Wagner et al., 1990). Their experimental results and analyses for C2H5 0 2 and C3H7 0 2 indicate that the alkene-forming process is not a separate reaction as indicated above. Instead it proceeds by a coupled path involving a second decomposition channel of the RO2* formed in which it first undergoes isomerization by internal H atom transfer:
+
-
+
+
+
R + O,-RO,*
+ M - RO, + M R02 + RH ROOH + R R02*
+
R02* QOOH*
- QOOH"
-
alkene
+ HO,
(5) (6) (7) (8) (9)
As shown, the excited RO2 formed in reaction 5 can decompose back to reactants, become stabilized by reaction with a third body, or isomerize to a hydroperoxyalkyl radical by intramolecular H atom transfer. As temperature increases, more and more of the RO2*goes on to form the alkene. Thus, reaction 4, as written, is really a global step. Nevertheless, the net result of both mechanisms above is the formation of the conjugate alkene t o the fuel. For the C2H5-02 system Baldwin et al. (1986), McAdam and Walker (1987), and Bozzelli and Dean (1990) presented alternative reaction channels for the decomposition to other products in addition to the alkene-forming path. On the basis of these results, early seminal work on alkylperoxy radical isomerization by Fish (19641, recent work on hydroperoxypropyVO2 reactions by Bozzelli and Rtz (19941, and the experimental results of the present study, we attempt to apply this new information about the lowto-intermediate temperature mechanism transition to the n-butane system. Reactions of the Butyl Radicals and Formation of Oxygenated Products. The character of the products of n-butane oxidation a t different temperatures is determined by the fate of the primary butyl (p-butyl) and secondary (s-butyl)radicals. At temperatures above approximately 900 K, these radicals undergo primarily
2292 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995
,f3 scission to smaller alkenes and alkyl radicals. Be-
tween about 650 and 900 K, there are high concentrations of the conjugate alkenes 1- and 2-butene. At temperatures below 650 K, hydrocarbon products are dominated by oxygenated species. The originating source of these compounds is the addition of molecular oxygen to the butyl radicals forming butylperoxy radicals:
+
s - C ~ H ~0,
s-C~H~O,
(11)
If we assume that H atom abstraction from the initial fuel is by mainly OH and HOz radicals, then based on the rate coefficients for these abstractions given by Pitz and Westbrook (19861, the concentration of s-butyl should be at least 2-3 times that of p-butyl a t 600 K. Since reactions 10 and 11 have the same rate parameters (A = 3 x lo1,; E = 0 (Benson, 1982)),the levels of s-C4Hg02 will exceed that of p-C4Hg02. Aside from the decomposition back to reactants, there are three primary consumption paths for these radicals. The relative importance of each of these paths is dependent on the fuel and oxidizer concentrations. The first path involves the bimolecular abstraction of a hydrogen atom by RO2 to form alkyl hydroperoxides:
s-C,HgO,
+ C4HIo - s - C ~ H ~ O O+HC4Hg
(13)
The initial fuel is a source of H atoms for these steps, which propagate the reaction by regenerating another alkyl radical for every alkylperoxy radical consumed. The activation energy for abstraction of a secondary H is approximately 17 kcaVmol and that for a primary H about 3 kcal/mol higher (Griffiths and Scott, 1987). These reactions are important in liquid phase oxidation. In the gas phase, mixtures which are fuel rich are conducive to these steps. Furthermore, these steps involving abstraction from the fuel are fairly slow, compared with, for example, abstraction of aldehydic H atoms. Aldehydes such as formaldehyde and acetaldehyde are important reaction intermediates in the low temperature oxidation of n-butane and are sources of readily abstractable H atoms. The general steps for an aldehyde (represented as RCHO) are shown below: p-C,HgO,
+ RCHO
+
P - C ~ H ~ O O HRCO
+
+
(14)
+
s - C ~ H ~ O ,RCHO --c s - C ~ H ~ O O HRCO (15) The C-H bond strength for an aldehydic H atom is approximately 9 kcdmol lower than that of a secondary H atoms in n-butane. For acetaldehyde, E14 = 10.5 kcal/ mol (Griffiths and Scott, 1987). Thus, aldehydes can be good hydrogen donors and contribute to peroxide production, but only a t a time in the overall reaction when their concentrations build up to significant levels. The alkyl hydroperoxides formed above undergo homolytic fission of the 0-0 bond:
+ OH s - C ~ H ~ O O H s - C ~ H ~+OOH
P - C ~ H ~ O O H p-C,HgO -C
(16) (17)
These chain branching type reactions have a fairly high energy barrier (-43 kcal/mol activation energy) but
accelerate the rate of the overall reaction by producing two highly reactive radicals: butoxy and OH. The second pathway for the consumption of ROZis by radical-radical disproportionation reactions of the type:
+ R0,-RO + RO + 0, RO, + RO - RO + RO
RO,
(18) (19)
These steps also lead to the formation of butoxy radicals. Although these reactions are not chain branching, they do lead to an overall acceleration of the rate of oxidation by replacing relatively unreactive ROz radicals with more reactive RO radicals. Benson (1982)has noted the importance of these type reactions in hydrocarbon oxidation systems. Other studies on n-butane oxidation (Baldwin and Walker, 1973; Carlier and Sochet, 1975) have demonstrated the importance of these type reactions specifically in the n-butane mechanism. The relative importance of this path increases at conditions where the radical levels are high, such as in undiluted fuel/oxidizer systems. The major products resulting from the two reaction paths discussed above are butoxy radicals. Most of the oxygenated hydrocarbon species observed experimentally a t the low temperature can be explained by the subsequent reactions of the alkoxy (RO) radicals, especially the butoxy radicals. The possible reactions of the two different butoxy radicals include both direct decomposition and bimolecular reactions each leading to a variety of oxygenated species:
- n-C3H7 + HCHO s - C ~ H ~OC2H5 + CH3CHO
p-C,HgO
(20) (21)
+ CH3 (22) p-C,HgO + 0, - C,H&HO + HO, (23) s - C ~ H ~+O0, CH3COCH2CH3 + HO, (24) p-C4Hg0+ 0, - l,2-C4H80+ HO, (25) s - C ~ H ~+O0 2 2,3-C4&0 + HO, (26) p-C4Hg0+ n-C4Hl0- 1-C4HgOH+ C4Hg (27) s-C4Hg0+ n-C4H,o - 2-C4HgOH + C4Hg (28) s - C ~ H ~ OCzH,CHO +
+
+
Direct decomposition of butoxy (reactions 20, 21, and 22) appears to be favored over the bimolecular H atom abstraction by 0 2 at these low temperatures, as seen by the high yields of formaldehyde and acetaldehyde relative to butyraldehyde and butanone observed in the experiment. Based on the rate parameters for RO decomposition compiled by Hucknall(19851, reaction 22 is favored slightly over reaction 21 at 650 K ( K z ~ / ~ ~ = zz 0.8). Reactions 27 and 28 appear not to be important even at this fuel-rich stoichiometry,as neither C4 alcohol was observed in the experiment. Reactions 25 and 26 are possible routes for the formation of the Cq oxiranes 1,2- and 2,3-C4H& from the butoxy radical. However, forming these particular species from 1- and 2-butoxy would involve a highly selective H atom abstraction. Thus, it is likely that these reactions will not compete very effectively with the direct decomposition of butoxy.
Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 2293
Butylperoxy Isomerization and Subsequent Reactions. Another reaction path is needed to explain the presence of the oxiranes in the products and, more importantly, the presence of the O-heterocyclic compound tetrahydrofuran (THF). The presence of THF is a clear indication that cyclization is occurring. Thus, the third and most important path for the reaction of C4H902 radicals is isomerization by the intramolecular abstraction and transfer of a hydrogen atom from one location in the molecule t o the outer oxygen of the peroxy group, forming hydroperoxybutyl radicals. This is represented by
RO,
-
QOOH
(29)
The alkylperoxy radical isomerization reaction scheme was first put forth in the early 1960s (Walsh, 1962;Fish, 1964). This scheme is believed to be a very important route for the consumption of ROz for higher carbon number fuels (C =- 3) and is believed to be dominant for fuels with carbon number greater than 5. Benson (1982) has shown that the rates for isomerization can be fast relative to the competing ROz reactions, and thus, this scheme represents an important additional reaction path. Each butylperoxy radical can undergo three possible isomerizations. The rate of each of these isomerizations is dependent on the the type of C-H bond broken t o obtain the internal H atom and the size (number of atoms) in the ring transition state. p-Butylperoxy can undergo 1,4s, 1,5s, and 1,6p transfers. sButylperoxy can undergo 1,4s, 1,4p, and 1,5p transfers. Estimates of the rate parameters for these can be found in Fish (1968b),Baker et al. (1975),Walker (1975), and Pollard (1977). For C4H902 we can compare the rate of the isomerization path (reaction 29) with the rate of the H abstraction path (reaction 131, assuming we start with For reaction 29 we also assume a an s - C ~ H ~ O radical. Z 1,4s k atom transfer. Thus for the ratio of rates, we have k~9/(k13[C4H101), using k29 = 1.26 x lo1, exp(-26.5/ R79 and k13 = 4 x 1OI2 exp(-l7/R79. For the [C4H101 conditions of the experiment and a t 650 K, this ratio is approximately 345. Thus, for butylperoxy, the rate of isomerization exceeds the abstraction rate. This effect is more pronounced a t higher temperatures. Two subsequent reaction paths of the hydroperoxybutyl radicals involve decomposition to products:
-
+ HO, QOOH - O-heterocycle + OH QOOH
alkene
(30)
(31)
There is strong support for reaction 30 being the primary mechanism by which the conjugate alkene is formed a t these temperatures instead of reaction 4 (Slagle et al., 198413; Bozzelli and Dean, 1990; Kaiser et al., 1990). For the hydroperoxybutyl radicals the reactions are
Table 3. Internal H Atom Transfers and 0-Heterocycle-FormingPathways for p and s-ButylperoxyRadicals radical
type of transfer
decomposition product (SOH)
CCCCOOH CCCCOOH CCCCOOH CCC(O0H)C CcC(O0H)C CCCC00H)C
1,4s 1,5s 1,6p
1,2-C4HeO 1,3-C4HsO tetrahydrofuran 1,2-C4HsO 2,3-C4HsO 1,3-C4H530
UP 1,4s L ~ P
s-CqHg02 radical. Cis- and trans-2-butene can only result from the s-hydroperoxybutyl radical following radical. Rate pa1,4s isomerization of the s-C~H~OZ rameters for the general reaction 30 have been compiled by Pollard (1977). Activation energies range from 18 to 25 kcal andA factors range from 1 x lo1, to 5 x 1013 s-l. More recent calculations based on the CzH502 system for the decompostion of CzH400H to CZH4 and HOz showed some disagreement in the calculated activation energies for this process: 29 kcal (Baldwin et al., 1986), 15.6 kcal (Bozzelli and Dean, 1990), and 17.9 kcal (Slagle et al., 198413). The other decomposition path involves elimination of OH followed by ring closure to form O-heterocycles.This path is necessary t o explain some of the observed oxygenated products from n-butane oxidation. The O-heterocycles consist of oxiranes, oxetanes, and tetrahydrofuran. Each butylperoxy radical will lead to distinct products based on the type of isomerization. For example, tetrahydrofuran can only result from the decomposition of a p-hydroperoxybutyl radical, which follows a 1,6p transfer, as the oxygen is needed on the end of the molecule for ring closure. On the other hand, considering only ROz isomerizatioddecomposition for the moment, 2,3-butene oxide can only result from the s-C~H~OZ radical. The different types of p-C4Hg02 and s-C~H~OZ isomerizations and resulting products for reaction 31 are summarized in Table 3. Another very important reaction path of the hydroperoxybutyl radicals is the further addition of 0 2 followed by isomerization and decomposition. This path leads to noncyclic oxygenated hydrocarbon products. In a recent study, Bozzelli and Pitz (1994) determined rate parameters and product channels for the reaction of 0 2 with hydroperoxypropyl radicals. They found that, below 700 K, addition of 0 2 to hydroperoxypropylresults in a stabilized adduct (at pressures above 1atm) which undergoes intramolecular H atom transfer followed by decomposition to products. Above 700 K, the reverse reaction is favored yielding 0 2 and hydroperoxypropyl which can react by other channels such as reactions 30 and 31. Low temperatures and high pressures will favor this path. We can specify analogous paths for the hydroperoxybutyl-02 reaction: 0 2 adds to the hydroperoxybutyl radical: QOOH
+ 0,
0,QOOH
(35)
Isomerization occurs as an H atom is abstracted and internally transferred to the outer oxygen of the peroxy group: s-C4HsOOH
-
2-C4H,
+ HO,
(34)
l-Butene can form from p-hydroperoxybutyl radical following a 1,4s isomerization of the p-C4H902 radical. It can also form from the s-hydroperoxybutyl radical following a more difficult 1,4p isomerization of the
0,QOOH
-
HOOQOOH
(36)
The hydrogen atom abstracted is assumed to be from the same carbon to which the hydroperoxide group is bonded. This H is favored because the C-H bond is approximately 4 kcal weaker than a secondary C-H bond due to the presence of the C-00 bond (Bozzelli
2294 Ind. Eng. Chem. Res., Vol. 34, No. 7, 1995 Table 4. Isomerization and Decomposition Products of p- and s-Hydroperoxybutyl Radicals CCC(0z)COOH CCC(O0H)COOH CCC(O0H)COOH CCC(OOH)C=O OH CCC(O0H)C-0 CCC(O)C=O + OH CCCOC-0 CzH&HO + CHO CC(0z)CCOOH CC(O0H)CCOOH CC(O0H)CCOOH CC(OOH)CC=O OH CC(O0H)CC-0 CC(O)CC-O + OH CC(O)CC-O CH3CHO + CHzCHO C(O2)CCCOOH C(O0H)CCcOOH OH C(O0H)CC~OOH C(O0H)CCC-0 C(0OH)CCC-0 C(O)CCC-O + OH CHzO + CzHiCHO C(O)CCC-O CCC(0OH)COz CCC(O0H)COOH CCC(OOH)COOH CCC-OCOOH OH
--- ---
+
4
+ +
+
CCC-OCOOH-CCC-OCO+OH
- -+
CCC=OCO CHzO C2HsC.O CC(Oz)C(OOH)C CC(OOH)C(OOH)C CC(OOH)C(OOH)C CC(OOH)C-OC OH CC(OOH)C=OC CC(O)C-OC + OH CC(O)C=OC CH3CHO CH3C0 C(Oz)CC(OOH)C C(OOH)CC(OOH)C C(OOH)CC(OOH)C C(O0H)CC-OC OH C(O0H)CC-OC C(6)CC=OC OH C(O)CC-OC CHzO cHzCOCH3 4
+
--
-- + -- +
+
+
-
carbonyls
+ 20H
-
+
+
--
+
+
and Pitz, 1994). This step is then followed by sequential 0-0 ruptures and /3 scission to final products:
HOOQOOH
spectroscopic measurements were made of cool flames, has shown that the luminescence is due to excited formaldehyde (Pipenberg and Pahnke, 1957; Downs et al., 1953. This is consistent with the products formed by this reaction sequence. At the low temperatures there is some contribution to the product distribution from the path: RO2 (RH or RCHO) ROOH (R or RCO), ROOH RO OH, and RO products. This first step depends on the fuel concentration and somewhat on the nature and size of the fuel molecule. As was shown earlier for n-butane and the conditions of the current experiments, this step is generally slow compared with the RO2 isomerization. However, the production of the C4 carbonyls butyraldehyde and butanone observed in the experiment are attributed to this path. Both are formed from the s - C ~ H ~radicals O (reactions 21 and 22). This is totally consistent with the findings of Sahetchian et al. (19901, that 90% of the total butyl hydroperoxides are of the secondary type. As temperature increases and approaches the intermediate temperature regime, RO2 isomerization becomes even more important relative to other RO2 paths. Moreover, the reverse reaction of hydroperoxybutyl 0 2 becomes favored allowing more QOOH to react to the O-heterocycles (reaction 31 and Table 3) and especially the butenes (reactions 30-32) and less to react to carbonyls (Table 4). The experimental results show a drop in the carbonyls and an increase in the butenes. This is consistent with other experimental findings (Salooja, 1967; Hucknall, 1985). As for the O-heterocycles, we see a slight decrease in the concentration of 2,3-C4H80 with increasing temperature, but this is difficult to assess since there are also low temperature paths forming it. We do see, though, from the experiment the appearance of tetrahydrofuran in the products a t the higher temperature which supports the mechanistic explanation. In addition t o the shift in the nature of the product distribution with temperature, the mechanism is capable of explaining the occurrence of the negative temperature coefficient. The negative temperature coefficient is a consequence of the shift in the important reaction paths in the mechanism in going from low to intermediate temperatures and the resulting changeover in the chain branching processes in each of those regimes. At low temperatures the reaction is accelerated by the production of two OH resulting from the addition of 0 2 to hydroperoxybutyl and its subsequent isomerization and decomposition. As temperature is increased to about 650 K, the reactants in QOOH 0 2 == O2QOOH become favored, which allows the competing paths QOOH butene + HO2 and QOOH O-heterocycle OH to take over. Thus, two OH radicals are replaced by OH and relatively unreactive HOz, causing a reduction in the overall rate of oxidation. Cool flames form in the low temperature regime by the low temperature mechanism. The resulting heat release raises the temperature of the system which causes the low temperature mechanism to be “switched off’. If there is sufficient fuel left, this can result in multiple cool flame oscillations, as seen in the experiments. In their modeling of a two-stage n-butane flame, Corre et al. (1992) found the O2QOOH to be the source of the OH prior to the cool flame and QOOH t o control the OH formation throughout the cool flame. The end of the NTC region is the temperature beyond which the rate begins to increase again. The HO2 produced with the mechanism shift leads to hydrogen
(37)
Since the 0-0 bonds are the weakest, two reactive OH radicals result from the process. This is significant because it makes this reaction path contribute substantially to chain branching. The details of the process represented above are specified for each of the individual hydroperoxybutyl radicals in Table 4. As seen in the table, the major carbonyl products resulting from these paths are formaldehyde and acetaldehyde. Baker et al. (1975) observed a strong dependence of the formation of aldehydes (CH20 and CH3CHO) on 0 2 concentration which supports their formation from this path. The paths specified in Table 4 are believed to be the primary reaction channels for hydroperoxybutyl plus 0 2 . Consideration can also be given to some alternative paths mentioned in the literature. In their study of the reaction of hydroperoxypropyl plus 0 2 , Bozzelli and Pitz (1994) found product channels which lead to olefinic hydroperoxide to be relatively important. These came about from the intramolecular abstraction of a hydrogen bonded to a carbon other than the one containing the OOH group. In another possible path, instead of isomerization, the hydroperoxybutyl-02 adduct abstracts a hydrogen intermolecularly to form a dihydroperoxide. Some of these compounds have been observed experimentally in the oxidation of longer chain hydrocarbons. Transition from Low to Intermediate Temperature Chemistry. The n-butane oxidation mechanism presented above is consistent with the measured species data from the experiment and the observed effect of temperature on the overall rate. At the lower temperatures the main path is 0 2 addition to butyl forming butylperoxy which isomerizes to hydroperoxybutyl radicals. These add with 02, isomerize, and decompose to 20H and aldehydes, mainly CH2O and CHsCHO, as observed. The two OH radicals produced greatly accelerate the reaction. This path can be responsible for the increase in the rate of oxidation which leads to the formation of cool flames. Previous work, in which
+
-
+
-
Ind. Eng. Chem. Res., Vol. 34, No. 7,1995 2295
750
-
0 Grlffitha at .I..
1993
0 Ramotowaki, 1892 0 hchrux del., 1971
both induction period and (dPldt),, criteria) and results from four other studies on n-butane in which the NTC region was determined. The mechanism and the model described by eq 38 predict reasonably well the onset of the NTC region especially for higher 0 2 (higher P)cases. The data of Nan-Chiang and Bardwell (1955) and Dechaux et al. (1971)were a t low pressure (-=0.15atm). The available reaction rate parameters used in eq 38 were near their high pressure limit. So it is reasonable that the agreement between the model and those data is not as good. Increasing pressure tends to shift the NTC region to higher temperatures. Increasing pressure will tend to preserve the low temperature mechanism as products will be favored in the QOOH 0 2 O2QOOH reaction. This will delay the onset of the NTC region. Importance of Secondary Reactions. Many of the stable products formed in the oxidation of n-butane are themselves highly reactive fuels, and thus become reaction intermediates. Chief among these are the aldehydes and alkenes. The aldehydes, especially acetaldehyde and formadehyde, are major sources of secondary reaction products. They are very reactive due to a readily abstractable hydrogen in the CHO group, and become the major sources of the carbon oxides. The alkenes, especially propene and the butenes, are also major sources of secondary products. This is particularly true in the intermediate temperature regime where the alkenes are formed early in the reaction and reach high concentrations. Subsequent reactions of the alkenes are the sources of the unsaturated aldehydes acrolein, methacrolein, and crotonaldehyde, as seen from the increased yields of these species a t 715 K (Table 2). They also can contribute to the formation of some of the saturated aldehydes, CO, COz, and lower alkanes and alkenes. An important submechanism in the oxidation of n-butane is the oxidation of methyl radicals. This is evident from some of the products formed in both temperature ranges. Methanol was the major hydrocarbon product formed in the experiment at 563 K. The methoxy radical is the source of methanol, forming by H abstraction from a hydrogen donor (fuel, aldehyde). Increasing the oxygen concentration will decrease the yield of methanol, and instead lead to more formaldehyde in the products, as the rate of the reaction of methoxy with 0 2 will increase: CH30 0 2 HCHO HOz. Methyl radicals obviously cannot form a conjugate alkene. Instead, as CH302 decomposes back to CH3 and 0 2 , the methyl will either abstract a hydrogen from the initial fuel or another molecule, leading to methane formation, or add with itself or another radical to form higher alkanes:
+
500
0
5
10
15
20
25
30
[02] x 10s (mollcma) Figure 8. Start of the NTC region for n-butane oxidation. Symbols represent experimental results. Solid curve represents best fit to data. Dashed curve represents calculated temperatures from eq 38.
peroxide. At temperatures within the NTC range, it is unreactive. At temperatures above approximately 700 K,depending on the pressure and other reaction conditions, hydrogen peroxide decomposes to OH: H202 + M OH OH M. Again there are two reactive OH radicals produced which accelerate the rate of oxidation and can lead to hot ignition. The mechanism discussed can be assessed t o determine if it can predict predict the temperature at which the NTC region begins. The onset of the NTC region should coincide with the changeover in the reaction path of the QOOH radicals. The competing steps are reactions 35 and 30. With respect t o the mechanism, we can define the temperature for the onset of the NTC as that value corresponding to some critical ratio of the rate of formation of alkene to that of 02QOOH:
-
+
+
d[buteneYdt - d[butenel d[OzQOOHYdt - d[O,QOOH] -
+
+
This equation was used to calculate the starting temperature of NTC behavior. The 1 atm data of Salooja (1967) for the onset of NTC behavior (642 K) was used to define the magnitude of the ratio in eq 38. The magnitude of eq 38 based on Salooja’s NTC starting temperature is 0.23. Using this magnitude, eq 38 was used t o calculate the NTC starting temperature over a range of initial oxygen concentrations. For reactions 30 and 35, the QOOH was assumed to be a p-CdHsOOH radical formed from a 1,4s H transfer. This would yield 1-butene, which was determined experimentally t o be the most abundant butene (Berry and Cullis, 1970; Berry et al., 1970). For the rate parameters, we used values from the recent work of Bozzelli and Pitz (1994). For reaction 30 we assumed A = 1.14 x 10l2s-l and E = 16.6 kcal/mol based on a similar reaction in the hydroperoxypropyl-02 system. For reaction 35 we used A = 3 x 10l2 cm3/(mol.s) and E = 0. In Figure 8 the computed results using eq 38 are compared with the current experimental results (using
CH3 CH3
+H
CH4
+ CH3 - C2He
-
-
(39)
(40)
All three of these products were measured in the experiment. Typically, small amounts of ethane and the next higher alkane from the initial fuel are formed in the products a t these temperatures. There are usually large concentrations of methyl radicals and fuel radicals (in this case butyl) present in an oxidizing alkane system. However, the abstraction path dominates at
2296 Ind. Eng. Chem. Res., Vol. 34,No. 7,1995
the pressure and temperatures of this study due to the large concentration of H atom donors. This is seen by the large yield of methane relative t o ethane and pentane in the measured products. Successful modeling of the chemistry of fuels such as n-butane require that these secondary reactions be considered and the correct submechanisms for the reactive intermediates be included in the mechanism.
Conclusions The oxidation of n-butane in the temperature range 550-740 K was studied. A region of negative temperature coefficient was observed experimentally between 640 and 695 K based on induction period data. Chemical analyses of the reaction intermediates and products revealed a shift in the oxidation mechanism across this temperature region. Oxygenated species dominate the hydrocarbon intermediates and products at low temperatures while alkanes and alkenes are the primary hydrocarbon species produced at intermediate temperatures. A mechanism was discussed in which the key low temperature steps are butylperoxy isomerization followed by 02 addition, further isomerization, and decomposition. In this scheme, two OH radicals are produced for every butylperoxy consumed and the main stable products are carbonyls. The two OH radicals accelerate the chain and lead to the formation of cool flames. As temperature is increased, the reverse reaction of QOOH 0 2 becomes important and the QOOH reacts through other channels, primarily to butenes and HO2 and to epoxides and OH. Also, direct decomposition of the butyl radicals begins to become important. The mechanism predicts a reduction in the rate of oxidation at temperatures consistent with the onset of the negative temperature coescient behavior observed in several different experiments. We are awaiting rate parameters and reaction channels, from measurements or calculations, of butylperoxy isomerization and subsequent addition with 0 2 , so we can perform detailed modeling of the low and intermediate temperature oxidation of n-butane.
+
Acknowledgment This work was supported by the U.S. Army Research Office under ARO Contract No. DAAG 29-85-K-0253, Reference No. 22437-EG.
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Benson, S. W. Cool Flames and Oxidation: Mechanism, Thermochemistry and Kinetics. Oxid. Commun. 1982,2, 169-188. Berry, T.; Cullis, C. F. Quantitative Studies of the Role of Alkenes During the Combustion of Alkanes. Aust. J. Chem. 1970,23, 2309-2314. Berry, T.; Cullis, C. F.; Trimm, D. L. Isotopic Tracer Studies of the Role of Butenes in the Combustion of n-Butane. Proc. R. SOC.London 1970, A316,377-389. Bozzelli, J . W.; Dean, A. M. Chemical Activation Analysis of the Reaction of CzHs with 0 2 . J.Phys. Chem. 1990,94,3313-3317. Bozzelli, J. W.; Pitz, W. J. The Reactions of Hydroperoxy-Propyl Radicals with Molecular Oxygen. Twenty Fifth Symposium (International) on Combustion;The Combustion Institute: Pittsburgh, 1994; in press. Bromly, J. H.; Barnes, F. J.; Mandyczewsky, R.; Edwards, T. J.; Haynes, B. S. An Experimental Investigation of the Mutually Sensitised Oxidation of Nitric Oxide and n-Butane. Twenty Fourth Symposium (International) on Combustion; The Combustion Institute: Pittsburgh, 1992; pp 899-906. Carlier, M.; Sochet, L. R. Detection by E.S.R. of Peroxy Radicals: The Importance of Radical-Radical Reactions in the Slow Oxidation of Butane. Combust. Flame 1976,25, 309-312. Carlier, M.; Corre, C.; Minetti, R.; Pauwels, J. F.; Ribaucour, M.; Sochet, L. R. Autoignition of Butane: A Burner and a Rapid Compression Machine Study. Twenty Third Symposium (Znternational) on Combustion; The Combustion Institute: Pittsburgh, 1990; pp 1753-1758. Cathonnet, M.; Boettner, J. C.; James, H. Experimental Study and Numerical Modeling of High Temperature Oxidation of Propane and n-Butane. Eighteenth Symposium (International) on Combustion;The Combustion Institute: Pittsburgh, 1981; pp 903913. Cernansky, N. P.; Green, R. M.; Pitz, W. J. Westbrook, C. K. Chemistry of Fuel Oxidation Preceding End-Gas Autoignition. Combust. Sci. Technol. 1986, 50, 3-25. Chakir, A,; Cathonnet, M.; Boettner, J. C.; Gaillard, F. Kinetic Study of n-Butane Oxidation. Combust. Sci. Technol. 1989,65, 207-230. Chandratna, M. R.; Griffiths, J. F. Pressure and Concentration Dependences of the Autoignition Temperature for Normal Butane Air Mixtures in a Closed Vessel. Combust. Flame 1994,99, 626-634. Cherneskey, M.; Bardwell, J. Surface Effefcts in Butane Oxidation. Can. J . Chem. 1960,38,482-492. Corre, C.; Dryer, F. L.; Pitz, W. J.; Westbrook, C. K. Two-Stage n-Butane Flame: A Comparison Between Experimental Measurements and Modeling Results. Twenty Fourth Symposium (International) on Combustion;The Combustion Institute: Pittsburgh, 1992; pp 843-835. Dechaux, J. C.; Flament, J. L.; Lucquin, M. Negative Temperature Coefficient in the Oxidation of Butane and Other Hydrocarbons. Combust. Flame, 1971, 17, 205-214. Downs, D.; Street, J. C.; Wheeler, R. W. Cool Flame Formation in a Motored Engine. Fuel 1953,32, 279-309. Dryer, F. L. The Phenomenology of Modeling Combustion Chemistry. In Fossil Fuel Combustion: A Source Book; Bartok, W., Sarofim, A. F., Eds.; Wiley: New York, 1991; Chapter 3, pp 121-213. Euker, C. A.; Leinroth, J. P. The Vapor-Phase Oxidation of n-Butane in a Flow Reactor. Combust. Flame 1970, 15, 275287. Fish, A. Radical Rearrangement in Gas Phase oxidation and Related Processes. &. Rev. (Chem. Soc. London) 1964,18,243269. Fish, A. The Cool Flames of Hydrocarbons. Angew. Chem. 1968a, 7, 45-55. Fish, A. Chain Propagation in the Oxidation of Alkyl Radicals. Oxidation of Organic Compounds, Gould, R. F., Ed.; Advances in Chemistry Series 76; American Chemical Society: Washington, DC, 1968b. Green, R. M.; Cernansky, N. P.; Pitz, W. J.; Westbrook, C. K. The Role of Low Temperature Chemistry in the Autoignition of n-Butane. SAE Pap. 1987, No. 872108. Griffiths, J. F.; Nimmo, W. Spontaneous Ignition Under Rapid Compression. Combust. Flame 1985, 60, 215-218. Griffiths, J . F.; Scott, S. K. Thermokinetic Interactions: Fundamentals of Spontaneous Ignition and Cool Flames. Prog. Energy Combust. Sci. 1987, 13, 161-197. Griffiths, J . F.; Coppersthwaite, D.; Phillips, C. H.; Westbrook, C. K.; Pitz W.J. Autoignition Temperatures of Binary Mixtures of
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Received for review September 12,1994 Revised manuscript received March 14,1995 Accepted April 10,1995@
IE940537J
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Abstract published in Advance A C S Abstracts, May 15,
1995.
