The chemistry of reactive radical intermediates in combustion and the atmosphere

The chemistry of reactive radical intermediates in combustion and the atmosphere CARRIGAN J. HAYES,a JOHN K. MERLEb and CHRISTOPHER M. HADADa a

Department of Chemistry, The Ohio State University, Columbus, OH 43210, USA Department of Chemistry, Winston-Salem State University, Winston-Salem, NC 27110, USA

b

1 2

Introduction 79 Basic concepts of combustion chemistry 80 Free radicals 80 Combustion at different temperatures 81 Methods for studying reactive combustion species 87 3 Reactive radical intermediates in combustion chemistry 91 Aliphatic systems 91 Aromatic systems 98 4 Future challenges in combustion chemistry 121 Fuel additives 122 Biodiesel 123 5 Conclusions 125 Acknowledgments 126 References 126

1

Introduction

Combustion processes convert chemical energy into heat and work, playing several important roles in today’s society. Oxidation processes provide power to beneficiaries ranging from automobiles to electrical generators; atmospheric oxidation reactions impact a wide range of environmental phenomena (i.e., ozone formation, photochemical smog, and acid rain). To fully understand combustion chemistry, it is necessary to understand the properties of the common reactive intermediates that participate in these reactions. Alkyl radicals (R•), alkoxy radicals (RO•), and peroxy radicals (ROO•) constitute the main classes of reactive radical intermediates involved in combustion. The occurrence and stability of the intermediates is governed by the temperature and pressure at which combustion (or oxidation) class takes place. Understanding these complex and dynamic relationships presents challenges for experimentalists and also for theorists. This review will first examine the fundamental chemistry that occurs during combustion of a fuel and then move into an exploration of some key classes of reactive intermediates. In particular, we will focus on the importance of reactive oxygen species (ROS) to combustion processes, highlighting our work on the unimolecular dissociative pathways available to the peroxy radicals of alkyl, aromatic, and heteroaromatic compounds. 79 ADVANCES IN PHYSICAL ORGANIC CHEMISTRY VOLUME 43 ISSN: 0065-3160 DOI: 10.1016/S0065-3160(08)00003-8

 2009 Elsevier Ltd. All rights reserved

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Basic concepts of combustion chemistry

FREE RADICALS

Any type of combustion chemistry is essentially dictated by the radical intermediates present. Before discussing how this occurs, it is first instructive to review some key aspects of radical chemistry. Most simply, free radicals are molecules with unpaired electrons. The lifetimes of these species vary widely given molecular composition and reaction environment, but the alkyl radicals of interest in combustion chemistry are highly reactive and thus short-lived. Radicals can vary in their hybridization (sp3, sp2, or sp), as well as the nature of the orbital in which their unpaired electron is placed (one dominated by s or p character); these orbital characteristics dictate the geometry around the radical center. Substituents (or functional groups) play a large role in stabilizing or destabilizing radicals. A free radical is stabilized by alkyl substituents on the radical center; tertiary radicals are more stable than secondary radicals, which are correspondingly more stable than primary radicals. Resonance stabilization (electron delocalization) also plays a role in radical stability: a radical adjacent to a p network (i.e., an allylic or benzylic system, or a carbonyl functionality) can delocalize the unpaired electron through this system and gain stability. Finally, inductive effects caused by an electronegative atom such as chlorine or a functionalized alkyl chain can affect radical stability. A radical is an odd-electron molecule, due to an unpaired electron, which can be oriented either spin-up (") or spin-down (#), leading to two degenerate electronic states as a doublet. The unpaired electron enables reactions that are different from those of closed-shell molecules. The basic steps of a radical chain reaction are familiar from undergraduate organic chemistry: initiation, whereby a reactive radical species forms; propagation, the processes by which the newly-formed radicals react with other molecular species to generate other radicals; and termination, via the collision of any two radical species to form one closed-shell molecule, thereby removing radicals from the system. Initiation can occur either thermolytically (when heat homolytically breaks a molecule’s bond) or photolytically (when high-energy light homolytically breaks a molecule’s bond). Once a radical forms, it can undergo a variety of propagation reactions, including hydrogen atom transfers, eliminations, additions, and unimolecular fragmentations; any reaction step that begins with a radical reactant and yields a radical product is classified as a chain-propagating step. Often, in flame chemistry, a reactive hydroxyl (HO•) radical is first formed, which then reacts with the fuel molecule via an initiation step (R—H þ HO• ! R• þ H2O). The ensuing variety of propagation possibilities constitute much of the chemistry of interest in combustion processes; this topic will be revisited shortly. Bond dissociation enthalpies (BDEs) aid in predicting relative reactivities for different organic (fuel) molecules. BDEs correspond to the enthalpy change for the homolytic cleavage of a chemical bond. The more stable the resulting radical is, relative to the reactant, the more favorable it is for the bond to break. General explanations for radical stability include alkyl substitution and resonance effects (as noted above). Additionally, Gronert has recently summarized work leading to an alternative explanation, that the

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potential for release of 1,3 repulsive energy (strain) has the greatest effect on these quantities: that is, a tertiary radical is more stable than a primary radical not because of the larger number of alkyl substituents attached to the radical center, but because of the greater magnitude of the geminal interactions in the parent (more congested) hydrocarbons that are relieved upon C—H bond cleavage to form the radical.1 Regardless of the origin of the effect, these BDE quantities are useful in rationalizing initiation steps for given fuels. Free radicals are involved in a wide variety of reactions, due to their reactivity and versatility; understanding their behavior as it relates to combustion is a goal of many experimentalists and theoreticians. Two specific classes of radicals of interest to combustion processes are peroxy (ROO•) radicals and oxy (RO•) radicals; along with hydroxyl (HO•) radical, many of these radicals are important members in the general class referred to as reactive oxygen species (ROS).2

COMBUSTION AT DIFFERENT TEMPERATURES

The concept behind combustion is straightforward – when a hydrocarbon fuel reacts with oxygen, the organic component is eventually converted to carbon dioxide and water – but the reality is more complicated. For instance, the combustion of methane (Reaction 1) is often used to teach students how to balance reaction equations: CH4 þ 2O2 ! CO2 þ 2H2O

ð1Þ

However, the combustion process for methane requires no fewer than 325 individual mechanistic steps (elementary reactions) to be accurately described, rather than the one-step route shown above.3 As such, incomplete combustion is a common occurrence and ROS are pervasive byproducts of that phenomenon, affecting an engine’s fuel efficiency and producing atmospherically detrimental emissions. Moreover, combustion varies with system temperature, as different oxidative pathways become accessible, as well as fuel/oxidizer ratio (equivalence ratio). By examining the representative cases of methane oxidation at high and low temperatures, this phenomenon becomes clearer. High-temperature combustion At high temperatures (as a general rule, T >1000 K), methane oxidation4 is initiated via hydrogen atom abstraction by hydroxyl radical, oxygen atom, or hydrogen atom (all of which are species generated in flames, oxygen to the smallest extent). Subsequently, in the most direct oxidative route (Fig. 1), methyl radical is oxidized to formaldehyde, which then loses a hydrogen atom to form formyl (HCO•) radical. Formyl radical can subsequently lose a hydrogen atom via collisional dissociation or reaction with molecular oxygen, thereby forming carbon monoxide (CO). A final oxidation step via the reaction of CO with hydroxyl radical yields the fully oxidized carbon dioxide as a final product. To more completely characterize the overall methane combustion mechanism,

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Fig. 1 High-temperature methane combustion. Adapted from Reference 7.

additional reactions such as methyl radical recombination and hydrocarbon eliminations, as well as ethene and acetylene oxidations, must also be considered. In this basic scenario, methyl, methoxy, and formyl radicals are intermediates of great interest; extending beyond this most straightforward case, the reactions of alkyl, alkenyl, alkynyl, alkoxy, and aldehydic radicals are all reactive intermediates of interest in high-temperature fuel combustion. Low-temperature combustion The low-temperature oxidation of methane (as a general rule, T < 1000 K) requires a more complex reaction scheme (Fig. 2). In reality, low-temperature oxidation of methane is unlikely to proceed readily, due to its substantial C—H bond strength; nevertheless, these pathways are shown below chiefly to illustrate common hydrocarbon oxidation pathways at low temperatures. The reaction scheme for oxidation of methane can be rationalized as proceeding in two phases.4 Since reactive flame species such as O(3P), H•, and HO• are not observed at low temperatures, the formation of the methyl radical must be achieved instead via an endothermic reaction with molecular oxygen.5 Once methyl radical is formed, it reacts with another oxygen molecule to form methylperoxy radical (CH3OO•), which abstracts a hydrogen atom from methane to form methyl hydroperoxide (CH3OOH) and methyl radical. Methyl hydroperoxide can unimolecularly dissociate to produce methoxy (CH3O•) and hydroxyl (HO•) radical. These last two steps are crucial, as they build up a reactive radical pool. Once a sufficient amount of methyl, methoxy, and hydroxyl radicals has formed, these species appropriate the initial duty of abstracting H atoms from methane, driving the reaction rate forward rapidly. Methyl radical is now formed at an appreciable rate and can undergo oxidation and recombination steps. More generally, low-temperature combustion relies heavily on the tendency of radical propagation to yield chain-branching reactions, a phenomenon first explored

THE CHEMISTRY OF REACTIVE RADICAL INTERMEDIATES O2

O2

•

•

CH3

CH4 •

•

CH4

83 •

CH3OOH + CH3

CH3O2

•

R = CH3, OH, CH3O • •

R•

CH4

CH3O + OH

•

CH3 + RH

•

•

CH3O2 •

O2

CH3OH + HCO

CH3OH •

CH3O + CH3O

•

CH3OH

CH3O2 •

CH3O2

•

CH2OH + CH3OH

CH2O

•

R

•

CH3OH + CH2O + O2

CH3OOH + HCO

O2 •

CH2O + HO2•

CH2OH + RH

• •

2CH3O + OH •

H + CO

•

HCO + RH

CH3OH + CH2O

Fig. 2 Low-temperature methane oxidation. Adapted from Reference 4.

by Semenov.6 Semenov’s reaction scheme is most relevant for species with two or more carbons and can be written as R• þ O2 ! alkene þ HO2•

ð2Þ

R• þ O2 þ M ! RO2• þ M •

ð3Þ

RO2 þ R—H ! RO2H þ R •

•

ð4Þ

•

ð5Þ

HO2• þ R—H ! H2O2 þ R•

ð6Þ

0

00

RO2 ! R CH(TO) þ R O

RO2H ! RO• þ •OH

ð7Þ •

R0 CH(TO) þ O2 ! R0 C (TO) þ HO2

•

ð8Þ

Using this scheme, we can track the original alkyl radical through the most likely mechanisms for oxidation at low temperatures that lead to chain-branching. Once formed from the parent molecule (R—H), an alkyl radical (R•) can react with molecular oxygen to form an alkene and hydroperoxyl (HO2•) radical [Equation (2)], via

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1,4-H-atom abstraction from the peroxy radical adduct. Alternatively, alkylperoxy radicals can be stabilized by collisionally transferring some excess energy, obtained during formation, to other molecules or atoms (M). The branching ratio between Reactions 2 and 3 is therefore highly pressure-dependent (i.e., it depends on the concentration of M). Once formed, the peroxy radical can abstract a hydrogen atom from an alkane to form a new radical [Equation (4)] or dissociate unimolecularly, yielding an aldehyde and alkoxy radical [Equation (5)]; both of these reactions produce a molecule that can participate in a chain-branching reaction. Since HO2• is unreactive at lower temperatures, Reaction 6 is less likely than the self-reaction of HO2• resulting in H2O2 and O2. Reactions 7 and 8 are chain-branching steps, which figure heavily in the exponential increase in radical concentration necessary to achieve ignition for the combustion of a given fuel (R—H). As shown, peroxy radical chemistry plays a substantial role in low-temperature combustion as opposed to the alkoxy radical chemistry of high-temperature combustion. Thus, the peroxy radicals constitute an important class of reactive intermediates with significant implications for low temperature combustion and atmospheric reactions. Negative temperature coefficient phenomenon In terms of temperature regions, low-temperature combustion occurs over the range 298–550 K, whereas high-temperature combustion mechanisms dominate at temperatures over 1000 K. Intermediate temperatures, from 550 to 700 K, demonstrate an unusual phenomenon called the negative temperature coefficient (NTC), which is observed for methane and larger hydrocarbon fuels.7 As shown in Fig. 3, when the correct alkylperoxy radical chemistry is included in a fuel’s combustion mechanism, a NTC range exists (Fig. 3, plot C) where an increase in temperature causes a decrease

Ignition delay

(a) (a) High T chemistry only (b, c) Peroxy radical chemistry included

(c)

(b)

Initial temperature

Fig. 3 Typical ignition delay of an alkane fuel as a function of the initial mixture’s temperature. Three different kinetic models are shown: (a) High temperature chemistry only; that is, no peroxy radical chemistry. (b) Same as (a), but the ‘‘Q•OOH’’ chain-branching channel of the peroxy radicals has been considered. (c) Same as (b), but the concerted elimination of RO2• to olefin þ HO2• has been considered (courtesy of Dr. Timothy Barckholtz, ExxonMobil Research and Engineering).8

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in reaction rate (i.e., a longer time to ignition).8 A simplified view9 of this phenomenon can be rationalized via an examination of the low-temperature combustion mechanism. Again, a crucial step in achieving the combustion of hydrocarbon fuels involves the unimolecular dissociation of an alkyl hydroperoxide (ROOH) into alkoxy (RO•) and hydroxyl (HO•) radicals, that is, chain-branching. The alkylperoxy radical leading to an alkyl hydroperoxide [Equation (4)] is in equilibrium with the reactants: alkyl radical and molecular oxygen (O2) [reverse of Equation (3)]. As the temperature increases, entropy favors the reactants, so that the alkylperoxy radical concentration will be minimized, and therefore, so will the alkyl hydroperoxide concentration. Thus, the chain-branching step cannot drive the oxidation forward, and the ignition time will increase. Moreover, for a hydrocarbon with two or more carbons, molecular oxygen can instead abstract a hydrogen atom from the alkyl radical, yielding an alkene and (notoriously unreactive) hydroperoxyl (HO2•) radical [Equation (2)]. It is also important to note the pressure-dependence of the competition between Equations (2) and (3); to persist, peroxy radicals must be stabilized after formation via collisions with other species. The NTC regime will remain in effect until the temperature increases sufficiently to allow for high-temperature, chain-branching pathways. At high temperatures, NTC is no longer relevant and the ignition rate increases with increasing temperature once again. Atmospheric oxidation The chemistry of the troposphere (the layer of the atmosphere closest to earth’s surface) overlaps with low-temperature combustion, as one would expect for an oxidative environment. Consequently, the concerns of atmospheric chemistry overlap with those of combustion chemistry. Monks recently published a tutorial review of radical chemistry in the troposphere.10 Atkinson and Arey have compiled a thorough database of atmospheric degradation reactions of volatile organic compounds (VOCs),11 while Atkinson et al. have generated a database of reactions for several reactive species with atmospheric implications.12 Also, Sandler et al. have contributed to the Jet Propulsion Laboratory’s extensive database of chemical kinetic and photochemical data.13 These reviews address reactions with atmospheric implications in far greater detail than is possible for the scope of this review. For our purposes, we can extend the low-temperature combustion reactions [Equations (4) and (5)], whereby peroxy radicals would have the capacity to react with prevalent atmospheric radicals, such as HO2•, NO•, NO2•, and NO3• (the latter three of which are collectively known as NOy): RCH2O2• þ NO• ! RCH2O• þ NO2• •

•

ð9Þ

RCH2O2 þ NO þ M ! RCH2ONO2 þ M

ð9aÞ

RCH2O2• þ NO2• þ M ! RCH2OONO2 þ M

ð10Þ

RCH2O2• þ HO2• ! RCH2OOH+O2

ð11Þ

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RCH2O2• þ HO2• ! RCHO þ H2O þ O2

ð11aÞ

RCH2O2• þ RCH2O2• ! RCH2O• þ RCH2O• þ O2

ð12Þ

RCH2O2• þ RCH2O2• ! RCHO þ RCH2OH þ O2

ð12aÞ

RCH2O2 þ NO3• ! RC(¼O)H þ HOx þ NOx

ð13Þ

Oxidation in the atmosphere begins photolytically with radiation from the sun rather than thermolytically; thus, atmospheric chemistry differs between day and night. In the daytime, the most common initiation step for VOC degradation involves photolysis of ozone by the sun’s ultraviolet light, leading to hydroxyl (HO•) radical generation: O3 þ h ! O(1D) þ O2 1

O( D) þ H2O ! 2HO

ð14Þ

•

ð15Þ

Once formed, the peroxy radicals have longer tropospheric lifetimes than HO• and can maintain larger concentrations. Common reactions of peroxy radicals include self-reactions to yield two carbonyl functionalized molecules and O2 [Equations (12) and (12a)]; reactions with HO2• to yield alkylhydroperoxides, aldehydes, H2O, and O2 [Equations (11) and (11a)]; and reactions with NOx species [NOx = NO• and NO2•] to yield alkoxy, alkylnitrate, and alkylperoxynitrate species [Equations (9), (9a) and (10)]. NOx reactions are of significant interest because alkylnitrate and alkylperoxynitrate molecules are stable enough to act as reservoirs, traveling long distances from VOC emission sources before decomposing, thus affecting distant air quality. For example, alkylnitrates can decompose to yield RO• and NO2•, and alkoxy radicals readily react with O2 to yield aldehydes. After H-atom transfer, the weak aldehydic CH bond is readily abstracted to yield an acyl radical, which can react with O2 and NO2• in succession to form peroxyacylnitrates (RC(¼O)OONO2 or PANs). PANs have various detrimental effects; in addition to the reservoir behavior described above, they are lachrymators and demonstrate mutagenic effects.14 Reactions of alkylperoxy radicals with NO3• can serve as an alternative nighttime source for HO• via a complicated series of reactions approximated by Equation (13). At night, when the sun’s radiation is minimal, the dominant VOC oxidant is nitrate radical (NO3•). The chemistry initiated by NO3• differs from that initiated by HO• radical in that NO3• prefers to react with unsaturated compounds via addition to one of the carbons of the p-system, rather than by hydrogen atom abstraction: NO2• þ O3 ! NO3• þ O2

ð16Þ

NO3• þ CH3CHTCH2 ! CH3•CHCH2ONO2 •

•

CH3 CHCH2ONO2 þ O2 ! CH3CH(OO )CH2ONO2

ð17Þ ð18Þ

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As shown, NO3• radical leads to different chemistry than does HO• radical; the peroxy radical can decompose to yield several products, including acetaldehyde, formaldehyde, 1,2-propanediol dinitrate (PDDN), nitroxyperoxypropyl nitrate (NPPN), and a-nitrooxyacetone. The reactions of the peroxy radicals with NOx species can lead to highly functionalized (and oxidized) organic compounds. The interplay of HO•, peroxy radicals, VOCs, and NOx species has substantial implications for tropospheric air quality. For instance, VOCs, NOx, and sunlight result in poor visibility from ozone and aerosol formation, together denoted as photochemical smog, which can lead to adverse health effects in sensitive individuals. Normally, we think of minimizing either class of compounds as beneficial to the atmosphere. However, minimizing VOC emissions only impacts ozone concentration in high-NOx areas. Moreover, in VOC-sensitive areas, reductions in NOx may lead to the overproduction of ozone. We can examine a simplified scheme15 for ozone production: NO2• þ h ! NO• þ O(3P)

ð19Þ

O(3P) þ O2 þ M ! O3 þ M

ð20Þ

In an ideal troposphere, O3 would react with NO• yielding NO2• and ultimately regenerating O3 [Equations (19) and (20)], thus no over-production of ozone could occur. However, in the presence of VOCs, the resultant peroxy radicals formed can compete with O3 by also reacting with NO• to form excess NO2•, thus resulting in the formation of excess ozone. This is just one example of the complexity of atmospheric chemistry; peroxy radicals and NOx have substantial implications for reaction with climate gases, acid rain formation, and other aspects of air quality. Saunders et al.16 and Jenkin et al.17 have provided a wealth of information on the tropospheric degradation of aliphatic and aromatic VOCs. Additionally, the interested reader may wish to consult References 10–17 for further discussion of these important topics. An additional concern in the commercial applications of combustion chemistry involves understanding and minimizing the production of harmful emissions [such as CO, carbon dioxide, and the oxides of sulfur (SOx), in addition to NOx] in combustion processes. Carbon dioxide increases the amount of greenhouse gas in the atmosphere and contributes to global warming, while NOx and SOx (oxidized derivatives of nitrogen and sulfur) can be transformed in water aerosols resulting in acid rain (via HNO3 and H2SO4, respectively). Another concern involves the formation of soot particles that can have severe respiratory effects. Many studies in combustion chemistry seek to reduce the production of these harmful byproducts.18

METHODS FOR STUDYING REACTIVE COMBUSTION SPECIES

Experimental methods Experiments performed in the 1960s and 1970s used a variety of approaches to detect molecular species and determine rate coefficients for the elementary steps that

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comprised hydrocarbon combustion. Many of these experimental methods are still widely used in combustion studies today. We will briefly discuss these methods as well as surveying some of the more sophisticated methods developed in recent decades. In combustion experiments, there are two key considerations: first, generating a flame and second, detecting the species of interest. Gaseous flows in a flame can be classified as laminar (streamlined layers) or turbulent. While these flames can be analyzed directly, it is less confounding to study flame chemistry through controlled generation of reactive species in one of a wide variety of experimental apparata. One such device is the shock tube.19 This cylindrical apparatus has both a highpressure region filled with an inert gas and a low-pressure region that contains the reactants of interest – fuel only if pyrolytic (thermal decomposition) processes are being examined, fuel and oxidizer if combustion processes are being examined – separated by a thin membrane. After the membrane is ruptured, a high-pressure shock wave travels down the low-pressure region and is reflected back on a microsecond time scale. The system temperature increases rapidly enough to yield hightemperature chemistry and is rapidly quenched for clean analysis. Other techniques simulate postignition flame processes. In the flow reactor, reactants of interest enter the reactor at one end and travel through a constant temperature region.7 Ideally, all concentrations and temperatures are consistent across the cross section of the reactor, so that all movement is in one direction and wall reactions are minimized. A similar approach involves the use of crossed molecular beams,20 wherein two molecular beams are directed into one another; the area of collisional intersection demonstrates chemistry that can occur in flames. Additionally, specific flame species of interest can be directly generated. Radicals can be generated photolytically (via light), thermolytically (via heat), chemically, or via microwaves. Laser-based methods are used to photolytically generate radicals. However, not every radical of interest can be generated from a convenient precursor; moreover, radicals generated via photolytic methods have excess internal energy, which increases the potential of their side reactions. Likewise, pyrolytic sources can be used with a wider range of species, but often necessitate a long residence time (ms) for radicals in the heating zone, giving high probabilities for radical–radical reactions and radical–wall reactions. These drawbacks have been circumvented via novel instrumentation in certain cases. For instance, Chen et al. developed a hyperthermal pyrolytic nozzle21 that works on a shorter time scale (ms) and generates thermally cold radicals via a jet expansion; this nozzle has been coupled with several targets and techniques by various researchers.22 In terms of coupling flame generation and detection methods, several combinations are common. Generally, shock tubes are coupled with IR and UV absorption and gas chromatography (GC) detectors, while flow reactors are used in tandem with GC, electron spin resonance, and resonance fluorescence detection. More advanced techniques are also available. For instance, mass spectrometry (MS) can be used by converting neutral radicals to ions, via chemical ionization and electron impact; these ions are then separated and detected according to their massto-charge ratios (m/z). The primary drawback of this method is that it cannot directly

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discern between isomers; structural isomers have identical m/z ratios but may behave in chemically unique fashions. Moreover, when MS is coupled with a flamegenerating system, the probe location must be chosen carefully, as the species population can vary greatly given the distance from the flame. Instrumental methods have become more sophisticated to face these challenges.23 In particular, Westmoreland and Cool have developed a flame-sampling mass spectrometer that has provided several revelations in terms of relevant molecular intermediates in combustion.24 Their setup couples a laminar flat-flame burner to a mass spectrometer. This burner can be moved along the axis of the molecular beam to obtain spatial and temporal profiles of common flame intermediates. By using a highly tunable synchrotron radiation source, isomeric information on selected mass peaks can be obtained. This experiment represents a huge step forward in the utility of MS in combustion studies: lack of isomer characterization had previously prevented a full accounting of the reaction species and pathways. When paired with an appropriate radical target, laser-based methods can serve either diagnostically, to discern which intermediates are present in a flame, or analytically, to explore the kinetics and dynamics of elementary steps of interest. Laser-induced fluorescence (LIF) and Raman spectroscopy (RS) are two typical diagnostic techniques; the former explores electronic transitions of a radical of interest, while the latter explores structural aspects of a species via observation of changes in its polarizability. Laser flash photolysis (LFP) is a common kinetic technique, in which a radical precursor is generated via a laser pulse, and the resultant radicals react with a target of interest; the rate of this reaction can be thus extrapolated by monitoring decreasing radical concentration over time. Many laser techniques are currently available and can provide a wealth of information on the structural, vibrational, and electronic properties of reactive radical intermediates. Several reviews of the use of lasers in combustion chemistry have been compiled, including those by Wolfrum,25 Crosley,26 and Eckbreth.27 Computational methods Computational and experimental methods clearly benefit from a symbiotic relationship in combustion studies.28 Theoretical calculations can propose important pathways to yield empirically observed intermediates by providing reaction energies and rate coefficients of elementary reactions, thereby guiding experiments. Moreover, theoretical calculations can potentially fill some gaps caused by limitations in experimental approaches: the vast majority of analytical techniques fail to distinguish between structural isomers and to identify short-lived intermediate species, both of which are important objectives in delineating overall combustion behavior. Finally, modeling can identify species to look for experimentally. Quantum chemical calculations are the most accurate theoretical methods available for studying the structures, energies, and elementary reactions of molecules. It is possible to determine the structure, energy, and geometrical parameters (i.e., vibrational frequencies, electronic states, and rotational constants) for reactants, transition states, and products of a chemical reaction. With this information,

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reaction thermodynamics [enthalpy (H), entropy (S), and free energy (G)] at specified temperatures can be estimated and reaction barriers and energies predicted. Furthermore, since reaction barriers can be calculated, absolute rate coefficients can be determined according to transition state theory (TST).29

k ðT Þ TST ¼ GðTÞ

kB T ðD G6¼ = kBT Þ 0 e h

ð21Þ

In Equation 21, T is the absolute temperature, h is Planck’s constant, kB is Boltzmann constant, and DG6¼ 0 is the free energy barrier height relative to infinitely-separated reactants. The temperature-dependent factor (T) represents quantum mechanical tunneling and the Wigner approximation30 to tunneling through an inverted parabolic barrier:   1 h i 2 ð22Þ GðTÞ ¼ 1 þ 24 kB T where  i is the imaginary vibrational frequency representing the curvature of the transition state barrier. Transition state theory yields rate coefficients at the high-pressure limit (i.e., statistical equilibrium). For reactions that are pressure-dependent, more sophisticated methods such as RRKM31 rate calculations coupled with master equation32 calculations (to estimate collisional energy transfer) allow for estimation of lowpressure rates. Rate coefficients obtained over a range of temperatures can be used to obtain two- and three-parameter Arrhenius expressions: Ea kðTÞ ¼ Aeð RT Þ

ð23Þ

Eo kðTÞ ¼ AT m eð RT Þ

ð24Þ

Common quantum mechanical methods for exploring the energetics of elementary reaction steps include ab initio33 and density functional theory (DFT).34 As computational speeds have increased, use of higher levels of theory have allowed for more accurate prediction of properties and reactions for reactive radical intermediates, further advancing our understanding of combustion chemistry. Using these methods, the elementary reaction steps that define a fuel’s overall combustion can be compiled, generating an overall combustion mechanism. Combustion simulation software, like CHEMKIN,35 takes as input a fuel’s combustion mechanism and other system parameters, along with a reactor model, and simulates a complex combustion environment (Fig. 4). For instance, one of CHEMKIN’s applications can simulate the behavior of a flame in a given fuel, providing a wealth of information about flame speed, key intermediates, and dominant reactions. Computational fluid dynamics7 can be combined with detailed chemical kinetic models to also be able to simulate turbulent flames and macroscopic combustion environments.

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Kinetic mechanism

Thermodynamic data

Reactor model

Results

Transport data

Fig. 4

Flow chart for a typical combustion simulation method. Adapted from Reference 35.

While theoretical calculations generally have been used to supplement experimental findings, they also hold enormous promise for fully discerning the potential energy surfaces of relevant combustion pathways, as well as identifying and exploring the chemistry of relevant reactive intermediates.

3

Reactive radical intermediate chemistry in combustion

ALIPHATIC SYSTEMS

Methane combustion The simplest hydrocarbon, methane, has posed a wealth of challenges to experimentalists and theoreticians seeking to discern its combustion mechanism. Methane’s reactions have been explored in a wide variety of contexts over the past several decades. We have discussed these briefly; the interested reader is referred to the reviews cited in our previous discussion for further details. Due to the scope of this review, we are primarily interested in these reactions insofar as they provide useful benchmarks for the reactions of larger alkylperoxy (RO2•) and alkoxy (RO•) systems. With respect to the reactive intermediates present in methane combustion and their implications for larger systems, Lightfoot has published a review on the atmospheric role of these species,36 while Wallington et al. have provided multiple overviews of gas-phase peroxy radical chemistry.37 Lesclaux has provided multiple reviews of developments in peroxy radical chemistry.38 Batt published a review of the gas-phase decomposition reactions available to the alkoxy radicals.39 Notably, the Gas Research Institute’s mechanism (GRI-MECH) for methane combustion3 is well-established, drawing on research from several groups over several decades to define and calibrate kinetic and thermodynamic data for each elementary reaction step. Additional mechanisms40 for methane oxidation are also available and updated periodically to include the most recent data. Methoxy (CH3O•) and methylperoxy (CH3O2•) radicals have been subjected to substantial study. Zaslonko et al. have reviewed several reactions involving the

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methoxy radical.41 More specifically, several experimental methods have been employed in studying this species. As early as 50 years ago, methoxy radical was an experimental target in mass spectrometric studies by Lossing.42 Burcat and Kudchadker used IR vibrational spectra of methanol at varying temperatures to extrapolate the ideal gas properties of methoxy radical.43 Ruscic and Berkowitz used photoionization MS44 to determine the ionization potential (IP) and heat of formation (DfH0) of methoxy radical. Martinez et al. used LIF to determine the temperature and pressure dependence of the rate coefficients for CH3O• þ NO2• over the 250–390 K temperature and 50–600 Torr pressure ranges.45 Wollenhaupt et al. examined the same reaction using pulsed laser photolysis (PLP).46 Computationally, Carter and Cook completed an extensive assessment of theoretical approaches for methoxy radical,47 while Page explored the kinetics of its unimolecular decomposition.48 Subsequent studies have focused on methoxy radical’s reactivity with a variety of molecules: Pan et al.49 and Sun et al.,50 with NO2•; Pang et al.,51 with NO•. Gomez et al.52 have modeled the addition and H-atom abstraction reactions of methoxy radical with various hydrocarbons. In early studies of methylperoxy radical, Simonaitis and Heicklin53 observed its reaction with NO• and NO2•, while Cox and Tyndall54 used molecular modulation spectroscopy to supplement these findings, and Kan and Calvert studied the water vapor dependence of the CH3O• self-reaction.55 More recently, Wallington et al. obtained absorption spectra of methylperoxy and other alkylperoxy radicals, over the 200–400 nm range.56 Tyndall et al. explored the self-reaction of methylperoxy radical via Fourier Transform Infrared (FTIR) spectroscopy, while Ghigo et al. have modeled the reaction energy surface computationally.57 Their findings complemented an earlier FP study by Lightfoot et al.58 Lesar et al. completed a quantum mechanical investigation of the reaction of CH3O2• with NO•.59 Biggs et al. explored the reaction of methylperoxy radical with NO3•60 as a possible source of nighttime HO• radical. Atkinson and Spillman have explored the kinetics of CH3O2• using cavity ring-down spectroscopy (CRDS), confirming previous findings by Hunziker61 and Pushkarsky.62 Enami et al. explored the kinetics of the reaction between bromine monoxide (BrO•) and methylperoxy radical.63 Together, these experimental and computational investigations provided mechanistic insights, rate constants, pressure dependences, branching ratios, and so on for the creation and validation of the 325step GRI-MECH for methane combustion. Ethyl radical þ O2 Intramolecular reactions become increasingly important as the size of the alkyl chain increases, for a given radical. Ethylperoxy (CH3CH2O2•) radical is a target of significant interest because it can undergo intramolecular rearrangement via a low-strain, fivemembered ring transition state structure. Historically, ethyl radical has been implicated as a discrete intermediate in many reactions. While early experimental studies of ethyl radical suggested that this species reacts bimolecularly with O2 to yield ethene (H2CTCH2, or C2H4) and hydroperoxyl (HO2•) radical,64 subsequent observation of other products such as acetaldehyde and oxirane imply a more complex mechanism. In

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1975, Hickel65 affirmed that this initial view of ethyl radical oxidation was incomplete, with support by Dechaux and Delfosse66 in 1979. Rather, it was deemed likely that ethyl radical oxidation proceeded through a Semenov-type mechanism [Equation (3)], in which ethylperoxy radical is an intermediate, formed by collisional cooling, which could transfer a hydrogen atom intramolecularly and decompose via multiple pathways. Subsequent experimental work bolstered the likelihood that ethylperoxy radical was an ethane oxidation intermediate. Baldwin postulated in 1986 that molecular oxygen adds to ethyl radical to form ethylperoxy radical, which then undergoes concerted H-atom transfer and elimination to yield ethene and hydroperoxyl radical.67 These and other experimental findings have been summarized in several reviews, including those of Fish,68 Walker,69 and Pilling et al.70 The small size of the CH3CH2• þ O2 system makes high-level computational exploration of its reaction energy surface tractable. Rienstra-Kiracofe et al. have provided an excellent review of the advances made in ab initio modeling of the CH3CH2• þ O2 potential energy surface over the past several years.71 These authors also summarized their own high-level calculations, noting the five most plausible pathways for the ethyl radical þ O2 reaction, four of which involve the unimolecular decomposition/rearrangement of ethylperoxy radical. C2H5• þ O2 ! C2H4 þ HO2•

ð25Þ

• • • C2H5• þ O2 !  CH3CH2OO !  CH2CH2OOH !  c-CH2CH2O+HO

ð26Þ

• • • C2H5• þ O2 !  CH3CH2OO !  CH3 CHOOH !  CH3CHO þ HO

ð27Þ

• • • C2H5• þ O2 !  CH3CH2OO !  CH2CH2OOH !  C2H4 þ HO2

ð28Þ

•

C2H5 þ O2 !  CH3CH2OO

•

• !  C2H4 þ HO2

ð29Þ

Equation (25) accounts for the NTC range observed for the ignition of ethane. Essentially, these reactions are refinements of the Semenov mechanism, since unimolecular reactions are important pathways in the oxidation of ethane. Due to its importance to hydrocarbon combustion as a model alkylperoxy radical, ethylperoxy radical continues to be the subject of experimental studies. Xing et al. used time-resolved, negative-ion MS to explore the reaction of ethylperoxy radical and nitric oxide.72 In separate works, Maricq and Szente explored the kinetics of ethylperoxy’s reaction with acetylperoxy radical73 and NO74 using transient diode laser absorption and time-resolved ultraviolet (TR-UV) spectroscopy. Atkinson and Hudgens have used ultraviolet cavity-ringdown spectroscopy (CRDS) to perform kinetic studies on the ethylperoxy radical self-reaction,75 while Rupper et al. used CRDS to explore the A˜ X˜ electronic transition of CH3CH2O2•.76 Hasson et al. used FT-IR and high-performance liquid chromatography (HPLC) with fluorescence detection to study the reaction of CH3CH2O2• with HO2•,77 while Mah et al. produced a mid-IR spectrum of this species.78 Ethoxy radical (CH3CH2O•) has enjoyed considerable interest as well. Choi et al. explored its photodissociation dynamics via photofragment translational

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spectroscopy,79 while Faulhaber et al. pursued the same goal using photofragment coincidence imaging.80 Computationally, Matus et al. performed CCSD(T) calculations81 to generate quantities of interest for this radical, including computing a heat of formation (–2.7  0.8 kcal/mol) via atomization energies. The development of an ethane combustion mechanism provides a historical context for understanding some overall trends of alkyl radical combustion. An understanding of the likely pathways for this small system is useful in modeling chemistry of larger systems, as can be observed from an examination of some other reactive radical intermediates. n-Propylperoxy radical (CH3 CH2 CH2 O2 ) In longer chain hydrocarbon radicals, isomerization reactions become more important; these pathways compete with bimolecular oxidation reactions and can impact ignition rates at low temperatures. For instance, n-propylperoxy radical can undergo several unimolecular dissociations/rearrangements:82 CH3CH2CH2OO• ! CH3CH2CH•2 þ O2

ð30Þ

CH3CH2CH2OO• ! CH3CHTCH2 þ HO•2

ð31Þ

CH3CH2CH2OO• ! CH3CH2CH•OOH

ð32Þ

CH3CH2CH2OO• ! CH3CH•CH2OOH

ð33Þ

CH3CH2CH2OO• ! •CH2CH2CH2OOH

ð34Þ

Reversion to reactants [Equation (30)] is straightforward; similarly, the concerted H-atom transfer and elimination pathway [Equation (31)] observed for ethylperoxy radical can also occur for CH3CH2CH2O2•. However, multiple isomerization pathways [Equations (32–34)] are now possible because the length of the alkyl chain has increased. In particular, the 1,5-H-atom transfer occurs via a six-membered ring transition state, thereby minimizing ring strain. Rearrangement reactions lead to a more complex mixture of products. Extrapolating from our knowledge of the ethylperoxy radical pathways, these isomerization products can decompose into propene, propanal, and other oxidation products. The preferences of n-propylperoxy radical will impact overall propane combustion. Because n-propylperoxy radical is small and can undergo low-barrier rearrangements, it is important to understand its possible rearrangement products; these findings have implications for larger hydrocarbons. DeSain et al. studied the production of hydroperoxyl radical via the reaction of n-propyl radical with O2 and proposed an activation barrier of 26.0 kcal/mol for Reaction 31;83 moreover, they have also observed sizable amounts of HO• radical, likely following one of the unimolecular H-atom transfer steps.84 Zalyubovsky et al. examined the A˜ X˜ • transition of CH3CH2CH2O2 (using CRDS) and detected several of its rotational isomers at 298 K.85b Chow et al. explored the kinetics of CH3CH2CH2O2• þ NO• via

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high-pressure chemical ionization mass spectrometry (CIMS).86 Kaiser explored the production of propene from the CH3CH2CH2• þ O2 reaction as a function of temperature and pressure and posited that chemically activated CH3CH2CH2O2• was an intermediate, resulting from the addition reaction’s exothermicity.87 Computationally, DeSain et al.88 performed QCISD(T) studies to explore the 1,4- and 1,5H-atom transfer reactions for several RO2• systems (including CH3CH2CH2O2•), generating RRKM/master equation rates to model HO2• and HO• production. Naik et al.89 performed a similar study using quantum RRK rates to yield rates for hydroperoxyl radical production in the n-propyl radical þ O2 system. Chen and Bozzelli have completed ab initio and DFT studies on the thermochemical and kinetic parameters of the n-propyl radical þ O2 reaction; they hypothesized that low-temperature ignition for this system results from intramolecular H-transfer reactions of CH3CH2CH2O2• followed by the addition of a second oxygen molecule making it capable of forming products that undergo chain-branching.90 Our group has examined the conformations of CH3CH2CH2O2•, as well as its possible unimolecular decomposition pathways.91 n-Propylperoxy radical conformations are described by two torsion angles: CC—CO in an anti orientation [here called trans (t)] or gauche (g) orientation and CC—OO in an anti orientation [trans (T)] or gauche (G) orientation. Energy profiles for these torsions were generated computationally, minima were obtained, enantiomers were noted, and five unique conformers were identified (Fig. 5). These conformers could rapidly interconvert, as the highest torsional barrier was predicted to be less than þ5.0 kcal/mol. Using the CBS-QB3 theoretical method,92 the 298 K distribution of rotamers was calculated to be 28.1, 26.4, 19.6, 14.0, and 11.9% for the gG, tG, gT, gG0 , and tT conformers, respectively. Therefore, all five conformers will be present as seen experimentally by Zalyubovsky et al.85b The unimolecular reactions of CH3CH2CH2O2• were studied in detail (Fig. 6); complete potential energy surfaces were generated using both DFT [B3LYP/ 6–31þG(d,p) and mPW1K/6–31þG(d,p)]34,93 and CBS-QB392 methods. As expected, 1,5-H transfer [Equation (34)] occurs with the lowest barrier, followed by simultaneous 1,4-H transfer and HO2• expulsion [Equation (31)]. The overall decompositions of each H-atom transfer product (i.e., each QOOH radical) were modeled. It

tG

gT

tT

gG′

gG

Fig. 5 Conformers of n-propylperoxy radical, named according to the conformational preferences around the central C—C bond, then around the C—O bond. [courtesy of John Merle (J Phys Chem A 2005;109:3637–3646). Reprinted with permission of J Phys Chem A.]

96

C.J. HAYES ET AL. ‡ 42.6 45.9 40.8

1,3-H transfer TS

‡ •

CH3CH2CH2 + O2 31.4 31.3 36.1

1,4-H transfer TS

Concerted 1,4-H transfer/elimination ‡

32.4 35.7 31.7 ‡

23.8 26.7 23.2

27.5 37.0 30.8

1,5-H transfer TS

ΔH(298 K), kcal/mol, relative to n-propylperoxy 0.0 B3LYP/6–31+G** 0.0 mPW1K/6–31+G** 0.0 CBS-QB3

Fig. 6 Initial reaction barriers for unimolecular reactions of n-propylperoxy radical. [courtesy of John Merle (J Phys Chem A 2005;109:3637–3646). Reprinted with permission of J Phys Chem A.]

was shown that, although the 1,5-H-transfer reaction has the lowest barrier, the resultant •CH2CH2CH2OOH cannot easily undergo further unimolecular rearrangements. Rather, the 1,4-H atom transfer routes [Equations (31) and (33)] encounter lower barriers in subsequent steps. This pathway for n-propylperoxy radical parallels a likely pathway involved in the decomposition of ethylperoxy radical.71 The CBS-QB3 potential energy surface accounts for the various experimentally observed products, including hydroperoxyl radical, propene, HO•, propanal, and oxirane (c-C3H6O). The activation barrier for simultaneous 1,4-H transfer and HO2• expulsion, obtained via calculations, compares well to the experimentally observed barrier (26.0 kcal/mol) of DeSain et al.83,84 This work provides some ramifications for larger alkylperoxy radicals: multiple conformers of long alkylperoxy radicals are likely to play a role in the overall oxidation chemistry and dictate consideration for correct treatment of thermochemistry; at lower temperatures (T < 500 K), unimolecular reactions dictate peroxy radical chemistry. n-Butoxy radical (CH3CH2CH2CH2O•) Just as n-propylperoxy radical is the smallest peroxy radical that can undergo the 1,5H-atom transfer, n-butoxy radical is the smallest alkoxy radical that can do so, while

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pentyl radical is the smallest alkyl radical with this potential. Similarly, these species are often used as models to understand the chemistry of larger alkoxy radical and alkyl radical systems. n-Butoxy radical exists in multiple conformer forms and can undergo a facile 1,5H-atom transfer. Vereecken and Peters94 exhaustively examined this possibility, via DFT calculations and TST rate coefficients, to recommend a rate coefficient of 1.4  105 s1, which agreed well with the experimental rates of both Atkinson95 and Hein et al.96 Vereecken and Peters used multiple approaches for deriving their multirotamer transition state theory expressions and demonstrated consistency through all of these. Ferenac et al. examined this unimolecular isomerization (1,5H shift) for n-butoxy radical and its functionalized derivatives, noting substantial substituent effects (more dependent on substitution patterns than on the functional groups themselves).97 Lendvay and Viskolcz examined unimolecular reactions available to n-butoxy radical via ab initio and RRKM calculations and noted that, while 1,5-isomerization was the fastest route, fragmentation reactions would compete at combustion temperatures.98 This finding was corroborated by exhaustive quantum chemical/RRKM dynamics calculations by Somnitz and Zellner.99 Jungkamp et al. generated an exhaustive atmospheric mechanism for n-butane via DFT and ab initio methods, proposing that n-butoxy radical will react primarily via 1,5-H transfers to ultimately form 4-hydroxy-1-butanal, while 2-butoxy radical will tend to decompose via b-scission to ethyl radical and acetaldehyde.100 Cassanelli et al. completed relative rate studies of 1-butoxy radical using FT-IR spectroscopy, noting that reaction with oxygen competed with isomerization.101 1-Pentyl radical (CH3 CH2 CH2 CH2 CH2 ) For pentyl radical, internal H-atom transfers can occur regardless of whether further oxidation occurs. These unimolecular reactions can directly compete with oxidation steps and so have implications for low-temperature combustion. For instance, n-pentyl radical can quickly isomerize to iso-pentyl radical via 1,4-H atom transfer; each of these radicals can undergo b-scission reactions to yield a new alkyl radical + alkene: C3 H7 þH2 C¼CH2

  n-pentyl !  iso-pentyl ! H2 C¼CHCH3 þC2 H5

ð35Þ

Several experimental studies, over the past several decades, have modeled the overall combustion of this system.102 Jitariu et al. have calculated unimolecular rates for reactions of pentyl radical (i.e., intramolecular H-atom transfer, b-scission, and elimination) noting that isomerizations have lower barriers than b-scissions.103 Larger aliphatic species In general, the trends predicted by n-propylperoxy radical, 1-butoxy radical, and n-pentyl radical provide good benchmarks for understanding the oxidation chemistry of longer chain radicals.104,105 For instance, 1,5-H-atom transfer and

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1,4-H-atom transfer concurrent with elimination provide two important routes with implications for both low- and high-temperature combustion. These possibilities constitute an important subclass of the reactions included in comprehensive mechanisms106 for the corresponding parent compound (i.e., propane,107 butane,108 and pentane). Also, developing larger mechanisms by substituting reaction data for similar, smaller species is a common practice in mechanism development (referred to as lumping109); thus, the reactions of these model compounds have parallels in the chemistry of larger fuels (in particular, n-heptane110 and iso-octane,111 which together constitute the primary reference fuels used to model gasoline combustion,112 as well as n-hexadecane, which shows promise for understanding diesel oxidation113). Chemistry taking place through five- and six-membered ring transition states is consistently favored kinetically over larger and smaller transition states. Gasoline and diesel fuels are two largely aliphatic hydrocarbon fuels that merit further discussion.114 These fuels are primarily used in different types of combustion environments: gasoline, in a spark-ignition (SI) engine, and diesel, in an autoignition engine. A SI engine relies on a four-stroke internal combustion process, involving the reciprocating piston, the intake valve, the exhaust valve, and a spark plug. In terms of the mechanism of combustion, a spark plug ignites the compressed fuel–air mixture, and the resultant flame ideally propagates smoothly across the engine cylinder. The speed at which the flame propagates is dependent on the fuel used. In a diesel engine, ignition relies on compression of the fuel until the autoignition temperature can be reached; no spark ignition is used, and no flame propagation occurs, and thus fewer emissions are involved. Similarly, autoignition temperature varies between given fuels. Even from this general overview, it is clear that the efficiency at which either type of engine operates depends heavily on fuel identity. Common metrics for understanding the ignitability of gasoline and diesel fuels are referred to as the octane number and the cetane number, respectively. Essentially, octane number refers to the volume percent of iso-octane in a given gasoline sample; cetane number refers to the volume percent of cetane (n-hexadecane) in a given diesel sample. In practice, the octane and cetane numbers refer to the practical efficiency of a given fuel to that of iso-octane and n-hexadecane, respectively. Wallington et al. have recently provided an excellent tutorial review of these and other related topics involving chemistry’s many roles in automotive fuels and engines.114

AROMATIC SYSTEMS

Soot formation Before we examine the oxidation pathways available to aromatic systems, it is first instructive to review the most notorious role of these compounds in combustion chemistry: their propensity to lead to soot formation. Soot is a byproduct of fuel-rich combustion, and soot particles can affect respiration and general health in humans.115 Soot production is a result of polycyclic aromatic hydrocarbon (PAH) formation in flames: as reactive hydrocarbon radical intermediates combine to grow

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and propagate, they can also cyclize into rings, which ultimately yield large networks of aromatic rings.116 To work toward minimizing harmful soot emissions, it is important to understand the mechanisms by which soot forms. Generally, the first cyclization step, whereby a benzene ring is formed from acyclic radical species, has been postulated to occur by one of two steps based on acetylene (C2H2) experiments: (1) the reaction of acetylene with either 1,3-butadien-1-yl radical (H2CTCH—CHTCH•) or buta-1-en-3-yn-1-yl radical (HCUC—CHTCH•) and (2) propargyl radicals (HCUC—CH2•) self-reaction, followed by H-atom transfers. More recently, cyclopentadiene has also been implicated as a likely precursor to benzene formation.117 Once the initial benzene ring has cyclized, it can undergo sequences of H-atom abstraction followed by acetylene addition, to yield PAHs. This is known as the H-abstraction-C2H2-addition (HACA) process, proposed by Frenklach and Wang.118 As an aromatic species aggregates to a size over 500 amu, it adopts a particulate form and can coalesce with other PAHs to further increase in size. When many of these particles agglomerate, they form soot.119 Efforts to minimize soot production are widespread. Notably, decreasing the carbon content relative to oxidizer concentration in a fuel/oxidizer mixture decreases the amount of soot formed.

Benzene and toluene In addition to their roles in soot formation, aromatic compounds undergo oxidation processes unique from acyclic saturated hydrocarbons. Aromatic species comprise 10–40% of gasoline and 5–30% of diesel; they reduce undesirable autoignition events (engine knock), thereby increasing a fuel’s octane rating. Therefore, aromatic oxidative decompositions have implications for combustion and atmospheric chemistry. The C—C and C—H bonds present in aromatic hydrocarbons are substantially stronger than those of alkanes, due to their sp2-hybridized carbons. When we consider the combustion reaction of a simple unsaturated species such as benzene or toluene, a new initiation step is possible. A reactive radical (for instance, HO•) may either abstract a hydrogen atom or add directly to the ring’s p-system, generating an allylic-type radical system: these multiple pathways compete (Fig. 7). Both pathways contribute to the combustion chemistry of aromatic species; HO• addition to the aromatic ring is the more prevalent at 298 K.120 In a monoalkyl-substituted aromatic species, such as toluene, abstraction of a hydrogen atom from the side chain can compete with HO• addition at positions ipso, ortho, meta, or para to the side chain. Thus, the possibilities for oxidative initiation and subsequent peroxy radical reactions increase.

Benzene oxidation Of the aromatic hydrocarbons, the oxidative pathways of benzene have been studied most exhaustively. Fujii et al.121 proposed a global mechanism in the early 1970s, in which the C—H bond of benzene is broken to form the phenyl (C6H5•) radical that

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C.J. HAYES ET AL.

Fig. 7 Reactions of toluene with HO• radical. HO• can abstract a benzylic hydrogen atom (a) or add to the aromatic ring at the ipso (b), ortho (c), meta (d), and para (e) positions relative to the methyl group. Each resultant radical can decompose by various pathways, depending on temperature and pressure.

subsequently reacts with molecular oxygen to form the phenylperoxy (C6H5OO•) radical: C6H5• þ O2 ! C6H5OO• ! 2CO þ C2H2 þ C2H3•

ð36Þ

Although greatly simplified, this model demonstrates the importance of phenylperoxy radical (C6H5OO•) as a reactive intermediate and accounts for some of the major combustion products. However, several other mechanistic intermediates (C3, C4, and C5 hydrocarbons) were also observed. Subsequently, Glassman’s mechanism for benzene combustion accounted for several more products, proposing that phenoxy (C6H5O•) radical was the chief reactive intermediate driving the combustion of benzene. The stepwise mechanism for benzene combustion122 began with H-atom loss to form a phenyl radical (C6H5•), then proceeded through reaction with O2, and then CO expulsion to form cyclopentadienyl (c-C5H5•) radical, which could react with O2 and expel CO again to form smaller hydrocarbon species (Fig. 8). However, this model still over-predicted the formation of phenyl, phenoxy, and cyclopentadienyl radicals; moreover, it failed to account for additional experimentally observed products (i.e., furan, pyranyl radical, and oxobutadiene).123 The tendency of the benzene combustion mechanism to substantially overestimate the formation of phenoxy radical suggested either flaws in the kinetic data involving C6H5O• or incompleteness of the combustion mechanism. Subsequent work verified the kinetics and energetics used in modeling C6H5O• decomposition. In separate studies, Liu et al.124 and Olivella et al.125 used ab initio and DFT models, along

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101

Fig. 8 High-temperature oxidation pathways of benzene, as proposed by Glassman in Reference 122.

with RRKM rate coefficients, to confirm the data for these pathways. (More recently, these data were also borne out in a quantum mechanical/RRKM study by Hodgson et al.126) Consequently, the benzene oxidation mechanism was further developed by considering additional decomposition and oxidation steps. Sethuraman et al. proposed that phenyl radical decomposition can occur by either of two key pathways:127 b-scission of phenyl radical or by breakdown of the phenylperoxy radical formed by the oxidation of phenyl radical (Fig. 9). Using PM3 calculations,128 which were ultimately verified by DFT studies,129 Carpenter predicted that another species, 2-oxepinoxy radical (3 in Fig. 9b), is an important intermediate due to its relative stability, formed via a spirodioxiranyl intermediate (2 in Fig. 9b) from phenylperoxy radical. Pathway A in Fig. 9b is the thermodynamically preferred pathway at temperatures increasing up to 432 K, while pathway B has an entropic benefit at higher temperatures. While pathway B essentially matched the traditional view of benzene combustion, pathway A introduced a new route for phenylperoxy radical, which could resolve discrepancies observed using previous models. This supposition was validated by experimental studies that demonstrated the prevalence of different ROS at different temperatures. Using CRDS, Yu and Lin studied the reaction of phenyl radical and oxygen, noting that phenylperoxy radical was the only adduct formed at temperatures ranging up to 473 K;130 Venkat et al. completed flow reactor studies of benzene combustion at 1200 K and identified phenoxy radical as a key intermediate.122 Our group has completed several studies of key reactions for some relevant reactive intermediates in benzene oxidation. Barckholtz et al. examined the oxidation pathways of several aromatic species, using benzene as a benchmark.129a General

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C.J. HAYES ET AL.

(a)

a b

– O

(b)

A

1

O•

B

O O

O• + O(3P)

• 2

6

O •

• O

O 7

3 –CO O

• O 4

• –CO2 5

Fig. 9 (a) Enthalpies (kcal/mol) leading to 2-oxepinoxy radical formation via unimolecular rearrangement of phenylperoxy radical. PM3/UHF = DfH and DFT (B3LYP/6–311 þ G(d,p)// B3LYP/6–31G(d)) = DH298. (b) Potential decomposition pathways for phenylperoxy radical A involving 2-oxepinoxy radical and B involving oxygen atom loss. [courtesy of Steven Kroner (J Am Chem Soc 2005;127:7466–7473). Reprinted with permission of J. Am. Chem. Soc.]

models were proposed for aromatic hydrocarbon and heterocycle oxidation (namely, benzene, pyridine, furan, and thiophene) in preparation for more specific studies of each relevant species. Carbon-centered radicals at each relevant position underwent exoergic oxidation, and the resulting peroxy radical unimolecular decomposition pathways were delineated. It was proposed that these peroxy radicals could undergo rearrangements that make them of significant atmospheric interest. A subsequent study examined phenylperoxy radical in greater detail. Fadden et al.129b identified five possible unimolecular decomposition pathways for phenylperoxy radical (Fig. 10): via oxygen atom loss to form phenoxy radical (Fig. 10, route A), via a dioxiranyl radical species (Fig. 10, route B), via a dioxetanyl radical

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103

Fig. 10 Potential unimolecular reaction pathways for phenylperoxy radical. Adapted from Reference 129b.

intermediate (Fig. 10, route C), via a 1,3-peroxy radical species (Fig. 10, route D), and via a p-phenylquinone radical intermediate (Fig. 10, route E). These routes were examined using DFT (B3LYP/6–31G(d)) and ab initio (CASSCF) structures, with high-level CAS-MP2, and UCCSD(T) single-point energies. Formation of the dioxetanyl species was predicted to be most favorable based on DG298 energies. This radical could unimolecularly decompose to produce smaller species, including cyclopentadienyl and pyranyl radicals, as well as acyclic oxygenated species, that are experimentally observed. At higher temperatures (T > 500 K), the entropic benefit afforded via oxygen atom loss becomes a contributing factor, and pathway A is the major decomposition route. Moreover, route B leads directly to 2-oxepinoxy radical (3 in Fig. 9b), which is potentially an important intermediate in low-temperature benzene oxidation. The stability of 2-oxepinoxy radical qualified it as a target for further theoretical and experimental study. The calculations of Barckholtz et al.129a allowed the refinement of a feasible energetic pathway toward 2-oxepinoxy radical; these DFT calculations supplemented the semiempirical work of Carpenter and also proposed a triradical intermediate between the dioxiranyl and oxepinoxy species (Fig. 9a). Consequently, the unimolecular decomposition of 2-oxepinoxy radical (3 in Fig. 9b) was thoroughly modeled by Fadden et al. using DFT (B3LYP) methods.131 Gibbs free energy profiles (T = 298–1250 K) were generated. A wide range of decomposition pathways were examined, which could account for typical experimentally observed products (Fig. 11). Notably, the delineated decomposition pathway did not require the generation of cyclopentadienyl radical. Cyclopentadienone is a commonly observed product in benzene combustion, and most mechanisms presume that cyclopentadienyl radical is its most likely precursor. However, it was shown that

O

40.8 + 2CO

H + CO 15 +30.3

16 +20.2

18 +54.2

13 +27.9

17 +26.2

49.0

+ CO H 19 –11.7

12.0

O CH2 C

H

O

3.4

6 +0.1

O

36.7

+ CO

12 +4.3

7 + 14.1

8 +36.6

H

O

C 6.8 36.3 O

+ HCO

C 22 +46.7

O C

O C

C

21 +25.0

+ CO 9 +39.0

O 30.5

C

H

2.7

6 +0.1

O

O

7.6

O + CO

O + CO

5 +28.9 C

H + CO

0.3

O

O

35.5

O

+ CO2 4 –28.1

3 +17.3

CO

23.4

7.7 H

2 +10.0

10.0

+ CO C 14 +3.3 O

–0.7

8.6

CH2

+ 2CO

CO2

39.4

43.8

29.4 4.6

O

8.8

+ 2CO

C2H3 + C2H2 + 2CO

104

O

O CO

O C

O C 26.6

34.7

13.3

2.8 H O

O + C2H2 + CO

40.2

11 +65.7

41.3 O

O H 30 +22.9

27.6

O O

C 29 +6.1

C 19.5

28 +7.3

40.6 C

26 +22.4

O

25 +42.3

O C

8.6

+ C2HO

C + C2H2O 27 +47.5

O 34.7

CO O 5 +28.9

0.3

O + CO 6 +0.1

Fig. 11 Unimolecular decomposition pathways of 2-oxepinoxy radical (1). The relative free energies (298 K, kcal/mol) at the B3LYP/ 6–311 þ G(d,p)//B3LYP/6–31G(d) level are shown for each intermediate relative to 1, and each free energy of activation is relative to the reactant for that specific step. [courtesy of Michael Fadden (J Phys Chem A 2000;104:8121–8130) Reprinted with permission of J Phys Chem A.]

C.J. HAYES ET AL.

+ C2HO 31 +34.2

22.8

O

23 +7.7

50.3 O

10 +38.8

O C O

20 +10.9 H + CO

C 24 +24.6 O

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2-cyclopenten-1-yl radical (19 in Fig. 11), a feasible source of cyclopentadienone, could be formed from the 2-oxepinoxy radical. These computed pathways supported the experimental results of Pfefferle and coworkers,140b in their work on the lowtemperature combustion of benzene, in which C2, C3, and C4 hydrocarbons were observed without formation of C5 constituents. 2-Oxepinoxy radical is highly stabilized by its unique structure; Mebel et al. determined, via PUMP3/6–31G(d)//UHF/6–31G(d) calculations, that it would have a relative enthalpy of DH0 = –91.8 kcal/mol as compared to phenylperoxy (C6H5O2•) radical.132 The stability of this species suggests that it will be fairly longlived and thus could undergo reactions of atmospheric interest, such as the addition of molecular oxygen. Merle and Hadad studied the oxygen-initiated decomposition of 2-oxepinoxy radical using both DFT and CBS-QB3 calculations.133 Calculations predicted that from T = 298–750 K, O2 addition routes compete with unimolecular decomposition (Fig. 12); for T > 750 K, however, the entropic penalty associated with the O2 addition step causes the unimolecular decomposition of 2-oxepinoxy radical to be more favored. The most stable O2 addition adduct below T = 1250 K is 6-peroxyoxepinone (route A in Fig. 12), which can cyclize to form a 1,4-peroxy intermediate which subsequently releases CO2 to form a 5-oxopentanalyl radical. This species can cyclize and fragment, yielding formyl radical, furan, and carbon dioxide. Above T = 1250 K, the

Fig. 12 Most favorable oxidative decomposition pathways for 2-oxepinoxy radical, at 298 K (path A) and 1250 K (path B). Adapted from Reference 131.

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dissociative pathway (route B in Fig. 12) afforded by 2-peroxyoxepinone’s loss of oxygen atom becomes more favorable. Even though 2-oxepinoxy radical (3 in Fig. 9b) may be an important intermediate in benzene oxidation, no confirmatory experimental evidence exists. Recently, Kroner et al. published an experimental study of the gas-phase acidity of 2(3H)oxepinone (C6H6O2), obtained via flowing-afterglow MS.134 They postulated that this quantity could be used along with a thermodynamic cycle to determine the heat of formation of 2-oxepinoxy radical: C6H6O2 ! C6H5O2 þ H þ

DHacid

ð37Þ

C6H5O2 ! C6H5O2• þ e

EA(C6H5O2)

ð38Þ

Hþ þ e ! H•

 IP(H)

C6H6O2 ! C6H5O2• þ H•

ð39Þ BDE=DHacid þ EA  IP

DHf(2-oxepinoxy radical, C6H5O2•)=BDE þ DHf(H•)  DHf(C6H6O2)

ð40Þ ð41Þ

A value of DHacid = 352  2 kcal/mol was determined for Equation (37). This experimental evidence could ultimately be valuable in conclusively identifying 2-oxepinoxy radical as a reactive species of interest. The overall pathways of benzene oxidation and the decompositions of possible intermediates have been well characterized via theoretical methods. Thus far, we have discussed these species mainly in the context of their oxidation mechanisms, but phenylperoxy and phenoxy radicals have also been investigated as individual experimental targets. The chemistry of phenoxy (C6H5O•) radical has been of interest for several decades. Benson et al. proposed the first rate coefficient for its unimolecular decomposition,135 while Lin and Lin provided information on the Arrhenius parameters for the reaction.136 Experimental and theoretical studies have examined phenoxy radical’s electronic states,137 molecular vibrational frequencies,138 and spin density, as well as its thermal and oxidative decomposition.139 Benzene combustion mechanisms rely heavily on the inclusion of phenoxy radical data;121,122,140 additionally, phenoxy radical decomposition reactions are necessary in the mechanisms for combustion of several other species, including propane,141 butane,142 and anisole.143 More notoriously, phenoxy radical has recently been implicated in routes to dioxin and polychlorinated naphthalenes; several theoretical studies have been completed on these potential reactions.144 Finally, the roles of phenoxy radical in flame chemistry,145 emissions,146 and other aspects of combustion147 have been compiled in several reviews. Phenylperoxy radical has similarly been a topic of experimental and theoretical interest. Tokmakov et al.148 calculated a potential energy surface for phenyl radical and O2 using ab initio G2(MP2) calculations. Weisman and Head-Gordon used timedependent density functional theory (TD-DFT) calculations to examine the effect of substituents on the phenylperoxy radical’s UV-vis absorption spectrum.149 Lin and Mebel used ab initio methods to study the phenoxy radical þ O-atom reaction.150

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Just et al. have examined the A˜ X˜ electronic transition of phenylperoxy using CRDS.151 Tonokura et al.152 used CRDS to study the visible absorption spectrum of the phenyl radical, as well as the kinetics of its reaction with O2. Krauss and Osman examined the UV absorption spectra of vinylperoxy radical (H2CTCHO2•) and phenylperoxy radicals.153 Phenylperoxy radical, originally assumed to be a factor in low-temperature combustion only, has actually been shown to play a substantial role in dictating the overall combustion trends of benzene. Just as the isomerizations and eliminations of the alkylperoxy radicals significantly affected their overall combustion pathways, rearrangements and other intramolecular pathways available to phenylperoxy radical similarly impact the overall progress of benzene combustion. This knowledge can be extrapolated to more complex aromatic species. Alkylated aromatics Like benzene, toluene (C6H5CH3) is a common constituent of gasoline. Much of the literature concerning toluene’s oxidation focuses on a global mechanism for understanding its combustion. Emdee et al.154 proposed that toluene’s combustion mechanism is most sensitive to its reaction with O2 to form benzyl radical (C6H5—CH2•) and HO2•. Dagaut et al. proposed that toluene oxidation is initiated by benzylic H-atom abstraction by O2 to form the benzyl radical, which can unimolecularly decompose to acetylene and cyclopentadienyl or react with an additional O2 and unimolecularly decompose to phenyl and formyl radicals (via benzaldehyde, (C6H5—C(TO)H).155 Pitz et al. generated a comprehensive mechanism for toluene combustion in varying settings, due to its widespread use as a fuel additive.156 El Bakali et al. noted an overall similarity between the chemistry of benzene and the chemistry of toluene, based on their oxidation mechanisms.157 Ethylbenzene has also been the subject of mechanistic studies,158 most recently by Ergut et al.159 As mentioned previously, the low-temperature oxidation of toluene is proposed to begin with either of two steps; if hydroxyl radical is present, HO• can abstract a benzylic hydrogen atom or add directly to the aromatic ring. Once a radical is generated on the aromatic ring or side chain of toluene, rapid oxidation can occur. The resulting peroxy radical has several viable pathways available. Atkinson compiled data on low-temperature atmospheric reactions for both the benzyl radical and the C6H5—CH3/HO• adduct.11 Andino et al. modeled the atmospheric oxidation of toluene and disubstituted xylenes, noting that several cyclized peroxy radical rearrangement products were energetically stable.160 Additionally, studies have focused on benzylperoxy radical (C6H5CH2O2•), although this is a less common target than phenylperoxy radical, because HO• addition to the aromatic ring is more dominant than H-atom abstraction from the benzylic C—H bond, in toluene combustion.120 Elmaimouni et al. studied the equilibrium for benzylperoxy radical and benzyl radical þ O2 over the temperature range 393–433 K, extrapolating the addition reaction enthalpy to be –20.1 kcal/mol at 298 K, and the free energy to be 11.4 kcal/mol.161 Fenter et al. performed a kinetic study162 using the same equilibrium at 760 torr with a temperature range of 298–398 K, proposing a

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rate expression kf(T) = (7.6  2.4)  1013 exp[(190  160)/T(K)] cm3/molecule-s, equilibrium Kp = 6.3  0.2  104 atm1, and reaction enthalpy DH298 = –21.8 kcal/mol. Buth et al. used a flow reactor coupled to mass spectrometric detection in a comparable study,163 determining kf = (4.44  1.3)  1011 cm3/molecule-s and Kp = 57,200 bar1 at 298 K. Noziere et al. monitored the reaction of benzylperoxy radical with hydroperoxyl radical using LFP/UV absorption and continuous photolysis/FTIR.164 El Dib et al. performed a LFP kinetics study of the self-reaction of benzylperoxy radical.165 It was again observed that rearrangement pathways comprise a substantial portion of the oxidation routes for alkylated aromatics.11,160 Since this phenomenon is mainly due to peroxy radical reactivity rather than to identity of the parent compound, it is clear that comparable rearrangements would be factors for PAHs, as well as for nitrogen-, oxygen-, and sulfur-containing heteroaromatic rings and their alkylated derivatives. Heteroaromatic combustion Considering additional functionalities in an aromatic ring allows for conclusions with implications for coal chemistry. Coal is a vital fossil fuel; about 50% of the United States is dependent on coal for electric power generation, and its use accounts for 90% of Ohio’s electrical power. Current clean-coal engineering efforts are underway to maximize coal’s energy potential while minimizing harmful environmental emissions (i.e., Hg, SOx, NOx, and CO2).166 Unlike hydrocarbon-based fuels like methane and gasoline, coal has never been subjected to a comprehensive mechanistic analysis, due to the complexity of its molecular structure. However, coal’s complex structure consists of various monocyclic units that can be explored: aromatic hydrocarbons and heteroaromatic rings are recurring units in coal’s structure, even while the overall structure varies geographically. Understanding low- and high-temperature oxidation reactions for these subunits and their reactive radical intermediates will facilitate a better understanding of their chemistry in combustion. Heteroaromatic compounds (Fig. 13) have been used as models for understanding coal chemistry in several pyrolytic studies. The azabenzenes (N-containing

Fig. 13 Heteroaromatic compounds of interest in coal combustion.

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heteroaromatics) have historically been most often studied; their pyrolysis has been shown to proceed through reactive radical intermediates although the identity of these intermediates can vary with reaction conditions.167 Using a shock tube, Mackie et al. observed the pyrolytic decomposition of pyridine,168 noting that three initial radicals derived from C—H bond scission: o-pyridinyl, m-pyridinyl, and p-pyridinyl. Of these, o-pyridinyl radical was dominant and was observed to yield cyanoacetylene (NUC—CUCH) through a ring-opening process, while m- and p-pyridinyl radicals formed HCN as a major product, via a less discernable pathway. Similarly, Kiefer et al. studied the pyrolysis of pyrazine, pyrimidine, and pyridine;169 all were observed to undergo ring-scission to yield 2-cyanovinyl radical (•CHTCH—CUN), which accounts for several combustion products upon decomposition, including HCN, NUC—CUCH, and acetylene. 2-Cyanovinyl radical has been identified in several pyrolytic studies of the azabenzenes, originally via the shock-tube study of Doughty et al.167f,g These authors proposed that the relative position of the nitrogen atoms in the ring substantially impacts the reactivity of a ring toward pyrolysis: for example, pyrazine dissociates more quickly than pyrimidine. Overall, the azabenzenes demonstrate increased reactivity relative to benzene; their C—H BDEs range from 93 to 98 kcal/mol, compared to benzene’s C—H BDE of 112 kcal/mol.170 The smallest C—H BDEs occur at positions ortho to nitrogen [although being a C—H bond that is twice ortho to N (as in pyrimidine) does not render a further lowering of the BDE].167 Kikuchi et al,171 Mackie et al.,172 and Jones et al.173 have attributed the lower BDE to the nitrogen atom’s in-plane lone pair electrons interacting with the unpaired electron on the carbon center of the radical, thereby reducing the strength of the C—H bond via radical product stabilization. The oxidative decomposition of the azabenzenes has not been studied in such great detail. Tabares et al. studied the reaction of pyridine with O-atom, noting a decrease in reactivity relative to benzene.174 Alfassi et al. studied the formation and reactivity of pyridylperoxy radicals in solution.175 Eisele postulated that the presence of ions derived from pyridine and picoline in the troposphere implicates these species as atmospherically significant.176 Yeung and Elrod explored this claim via chemical ionization MS to study the reactions of HO• with pyridine, the picolines, the lutidines, and the ethylpyridines and postulated that pyridinated compounds could indeed have substantial implications on tropospheric ion content.177 As with toluene, the reaction of HO• with pyridine and alkylated pyridines is likely to proceed either via HO• addition to the aromatic ring or hydrogen atom abstraction from a C—H bond. The new radicals can unimolecularly decompose or undergo reaction with O2. These reactions have been modeled for toluene and the hydrocarbon analogues by Andino et al.160 Heteroaromatic compounds have the potential to add O2 at the ring nitrogen and thus form NOx species,178 potentially leading to excess tropospheric ozone and acid rain.179 The five-membered heteroaromatic (furan, oxazole, pyrrole, and thiophene) are of additional interest. Besides their role in coal combustion, these have been implicated as emissions of biomass burning,180 residential fires,181 waste tire burning,182 cigarette smoking,183 and motor vehicles.184 Bruinsma et al. examined the pyrolysis of heteroaromatic rings most commonly found in coal volatiles,185 determining a rank of increasing pyrolytic reactivity: thiophene < benzene < pyridine < pyrrole < cyclopentadiene < furan; they also noted that an additional, fused aromatic ring

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has a stabilizing effect, especially for pyridine and furan. Cullis and Norris also studied the pyrolytic processes of the heteroaromatics, yielding methane and benzene as major products, via hydrocarbon radical intermediates. Qualitative product analysis revealed that these heteroaromatics decomposed by similar mechanisms regardless of identity; the heteroatoms were generally not present as major combustion products; they were lost as water, hydrogen sulfide, or hydrogen cyanide.186 Braslavsky and Heicklen extensively reviewed the thermal and photochemical decomposition of heteroaromatic compounds.187 Klein et al. studied variations in heteroaromatic C—H BDEs for substituted aromatic compounds.188 In particular, furan and pyrrole, as well as their methylated derivatives, have been common targets of pyrolytic studies. The pyrolysis reactions of furan and its methylated derivatives have been shown to lead to various products, including CO, acetylene, acetaldehyde, propyne, and allene. Grela et al. observed that methylated furan is likely to undergo C—O bond scission, yielding either benzene and water, via the biradical •C(CH3)TCH—CHTC(CH3)O•, or to isomerize prior to decomposition to produce CO and C5H8.189 Organ and Mackie also suggested that the biradical intermediate was the most likely intermediate.190 The mechanism has since been re-evaluated. Fulle et al. noted that the major decomposition products of unsubstituted furan were formed via one of two pathways,191 one which resulted in C2H2 and ketene (H2CTCTO) and one which led to propyne and CO; Sendt et al. confirmed these pathways and proposed that they were achieved via 1,2-H transfer in the original furan molecule, which led to cyclic carbene intermediates.192 The pyrolysis of pyrrole produces a variety of products: hydrogen cyanide, propyne, allene, acetylene, cis-crotonitrile, and allyl cyanide, among them. Lifshitz et al. hypothesized that pyrrole undergoes 1,2-bond (N—C) cleavage, then an internal H-atom transfer, to yield a radical intermediate that can isomerize to either cis-crotonitrile or allyl cyanide, or dissociate to HCN and propyne.193 Bacskay et al. completed quantum chemical comparisons of the isoelectronic pyrrolyl and cyclopentadienyl radicals; they hypothesized that pyrrolyl radical is formed via C—H bond scission in the intermediate pyrrolenine (2H-pyrrole) rather than directly via N—H bond cleavage (Fig. 14).194 Mackie et al. explained a similar finding, postulating that it was the formation of pyrrolenine that dictated the rate at which pyrrole pyrolysis occurred.195 While most studies have focused on the pyrolytic unimolecular decomposition of these monoheteroaromatic compounds, our group has explored their oxidative decomposition. As with benzene, where phenylperoxy radical plays a major role in dictating oxidation pathways, we hypothesize that the peroxy radicals derived from heteroaromatic rings are reactive species of considerable interest for combustion and atmospheric reactions.

Fig. 14 Likely pyrolysis pathway of pyrrole, via intermediate pyrrolenine. Adapted from Reference 194.

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Barckholtz et al. surveyed a variety of computational methods and basis sets to select an appropriate theoretical model for study of these molecules, finding that DFT provided a good balance of accuracy and computational economy.196 This study also validated the use of heteroaromatic monocyclic rings as constructive models for their polycyclic analogues (by extension, this finding could also confirm the similarities between the chemistry of coal and these smaller model compounds). BDEs were compiled for a variety of polyheteroaromatic rings; the resulting values were compared to the corresponding monoheteroaromatics. For example, for benzofuran, the C—H BDE at the 2-position of the furan ring is 117.8 kcal/mol, which is 0.6 kcal/mol less than the corresponding BDE in furan itself. These calculations showed that increasing the number of rings in the compound did not have a substantial effect on BDE, except in the case of a C—H bond adjacent to a bridgehead junction when an electronegative heteroatom was present on the other side of the bridgehead. Even in such cases, the deviation between the monocyclic analogue and the polycyclic derivative was only 2 kcal/mol. In separate DFT studies, Fadden et al. examined the rearrangement pathways (Fig. 15) of peroxy radicals from azabenzenes197 and five-membered heteroaromatic rings.198 It was observed that each azaphenylperoxy radical can lose molecular oxygen (2 ! 1), rearrange to a dioxiranyl species (2 ! 3) or a dioxetanyl species (2 ! 4), or lose atomic oxygen (2 ! 5). Other unimolecular decomposition pathways afforded to alkylperoxy radicals (i.e., H-atom transfer and b-scission) are not possible for their aromatic analogs. From the calculated energies, several main conclusions were drawn. Loss of O2 is less endoergic at 298 K than loss of O-atom for most heteroaromatic peroxy radicals; two exceptions are 3-pyridinylperoxy radical, which mimics phenylperoxy radical in this aspect of its reactivity as well as many

Fig. 15 Unimolecular pathways available to heteroaromatic peroxy radicals; example shown for 2-pyridinylperoxy radical. Adapted from Reference 198.

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others, and 3-pyridazinylperoxy radical, in which the oxy radical is a -radical (2A0 ) and maintains aromaticity not observed in the peroxy radical. Dioxiranyl formation is generally less endoergic than O2 and O-atom loss at 298 K; however, at temperatures greater than 500 K, the entropic contributions reverse the stabilities. Dioxetanyl intermediates are more strained and, therefore, unstable intermediates; a few exceptions are observed for compounds in which an alkyl chain is adjacent to a ring nitrogen and cyclization can form stable nitrosyl radicals (as for Pathway 4b for 2-pyridinylperoxy radical), but these reactions still incur high reaction barriers. Similar routes are available to peroxy radicals of O-, S-, N-, and O,N-containing five-membered ring heteroaromatics. The kinetic and thermodynamic parameters for the viable arylperoxy radical unimolecular dissociation steps were calculated for the furanylperoxy radicals; the effect of a second heteroatom was examined by a comparable approach for the oxazolylperoxy radicals. Thermodynamic parameters for these reactions were compared to those of the pyrrolylperoxy and thiophenylperoxy radicals. For smaller heteroaromatic peroxy radicals, loss of O-atom to form the corresponding aryloxy radical is preferred at 298 K to other decomposition routes. Dioxiranyl formation (Fig. 15, pathway 3) competes thermodynamically with oxygen atom loss (Fig. 15, pathway 5) in some cases and is universally more favorable than O2 loss (Fig. 15, pathway 1). The dioxetanyl route (Fig. 15, pathway 4) is disfavored for five-membered ring heteroaromatic peroxy radicals, often due to formation of an anti-Bredt double bond in the ring system. As with the azabenzenes, the dissociative pathways become more favorable than rearrangements at high temperatures. Overall, reactivity of the heteroaromatic peroxy radicals was shown to depend heavily on ring size. For azabenzylperoxy radicals, losing an oxygen atom is substantially unfavorable, and reversion to reactants (aryl radical þ O2) is a more likely dissociation pathway; the peroxy radical of the five-membered ring heteroaromatics can lose oxygen atom at a lower cost. For both sets of arylperoxy radicals, isomerization pathways are important at low temperatures. Additionally, intramolecular cyclizations compete with O-atom loss, and some cyclizations lead to nitroso radicals, creating implications via possible NOx formation. As with phenylperoxy radical, the azabenzylperoxy radicals have lower barriers for rearrangement than for loss of oxygen atom, and consequent products will influence overall combustion pathways. Alkylated heteroaromatics Substantially fewer studies have been published for the reactions of alkyl-substituted heteroaromatics, although these compounds also have implications for coal combustion. Several references discussed in the previous section contain information on methylated heteroaromatic rings. Mackie and coworkers completed experimental199,200 and theoretical201 studies of the pyrolytic decomposition of 2-picoline (2-methylpyridine). They concluded that decomposition proceeded mainly through o-pyridinyl and 2-picolinyl radicals. The former tended to decompose predominantly to yield cyanoacetylene, while the latter favored decomposition to a cyano-functionalized cyclopentadiene (Fig. 16).

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Fig. 16 Cyanocyclopentadiene.

Despite the authors’ assertion that alkylated heteroaromatic compounds provide a better model for fuel-bound nitrogen than do the unsubstituted heterocycles, their pyrolytic study remains the most comprehensive look at substituted heteroaromatic chemistry, even several years later.196–202 Kinetic studies are more common in the literature: Frerichs et al. examined the reaction of the picolines with oxygen atom,202 while Yeung and Elrod studied reactions of HO• with pyridine and its methyl- and ethyl-substituted derivatives.177 Both groups noted that the presence of nitrogen did not demonstrably affect the species’ chemistry; generally, reactivity is comparable to toluene. The oxidation pathways for alkylated heteroaromatics start with the formation of a radical species, via hydrogen atom loss or alkyl group homolytic bond cleavage. We calculated these BDEs for methyl- and ethyl-substituted derivatives of several key heteroaromatics (Tables 1–3).203 Few of these experimental values exist;204 therefore, Table 1 Thermodynamic and spin density information for methyl hydrogen atom loss reactions of methyl-substituted heteroaromatic rings. Enthalpies and energies in kcal/mol, obtained at the B3LYP/6–311þG**//B3LYP/6–31G* (designated as B3LYP) and CBS-QB3 (designated as CBS) levels

DH298

Toluene Pyrrole Furan Thiophene Oxazole Pyridine Pyridazine

DG298

( – )

Methyl

B3LYP

CBS

Experimental BDE

B3LYP

CBS

1 2 3 2 3 2 3 2 4 5 2 3 4 3 4

86.7 83.1 86.8 83.1 87.4 85.2 87.0 86.4 88.1 84.5 88.2 87.0 87.9 88.9 87.6

90.6 86.1 90.1 86.3 90.5 86.5 89.9 89.9 91.1 87.9 92.0 91.0 91.6 93.3 91.7

88.0–90.3a

79.4 75.3 78.9 75.3 79.5 84.2 88.1 78.7 80.2 76.6 80.5 79.6 80.6 81.3 80.5

83.8 78.9 82.4 78.6 82.7 79.1 82.2 82.3 83.3 80.1 84.7 83.7 84.6 85.8 84.6

96.0b

0.72 0.63 0.74 0.60 0.73 0.60 0.71 0.64 0.63 0.72 0.73 0.72 0.74 0.76 0.75

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Table 1 (continued ) DH298

Pyrimidine Pyrazine

Methyl

B3LYP

CBS

2 4 5 2

89.5 89.3 87.5 88.0

93.1 92.7 91.4 92.3

DG298 Experimental BDE

( – )

B3LYP

CBS

82.4 81.9 80.7 80.5

86.7 85.5 84.7 85.0

0.73 0.75 0.72 0.72

See Fig. 13 for structures and numbering. a Reference 158. b Reference 173.

Table 2 Thermodynamic and spin density information for ethyl (or methylene) hydrogen atom loss reactions of ethyl-substituted heteroaromatic rings; enthalpies and energies in kcal/ mol, obtained at the B3LYP/6–311þG**//B3LYP/6–31G* (designated as B3LYP) and CBSQB3 (designated as CBS) levels

DH298

Ethylbenzene Pyrrole Furan Thiophene Oxazole Pyridine Pyridazine Pyrimidine Pyrazine a

Reference 118.

DG298

( – )

Ethyl

B3LYP

CBS

Experimental BDE

B3LYP

CBS

1 2 3 2 3 2 3 2 4 5 2 3 4 3 4 2 4 5 2

83.9 80.8 83.8 80.1 84.2 79.8 83.5 82.4 84.5 81.5 84.2 82.9 83.7 84.7 83.7 84.2 83.9 84.3 83.8

88.1 84.6 87.7 83.9 87.9 83.8 87.1 86.5 88.3 85.6 88.1 85.6 87.2 90.1 88.7 89.0 88.5 89.1 90.2

85.4–86.9a

75.2 72.2 75.3 71.6 75.4 71.1 74.4 74.1 76.0 73.0 75.2 74.8 74.5 77.6 75.1 76.1 75.7 74.9 75.1

79.7 75.9 78.9 75.5 79.1 75.2 78.1 78.4 79.9 77.1 80.0 77.2 78.7 83.1 79.7 80.9 80.4 80.5 81.4

0.68 0.61 0.71 0.57 0.69 0.55 0.67 0.61 0.68 0.60 0.67 0.69 0.69 0.68 0.68 0.69 0.70 0.66 0.68

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Table 3 Thermodynamic information (kcal/mol, 298 K, B3LYP/6–311þG**//B3LYP/ 6–31G*) for alkyl group loss reactions of methyl- and ethyl-substituted heteroaromatic rings

Pyrrole Furan Thiophene Oxazole Pyridine Pyridazine Pyrimidine Pyrazine

Methyl

DH298

DG298

Ethyl

DH298

DG298

2 3 2 3 2 3 2 4 5 2 3 4 3 4 2 4 5 2

106.1 104.3 108.8 105.7 103.9 100.9 109.8 106.9 110.8 94.1 98.3 97.6 96.2 97.1 97.4 90.3 99.7 94.9

89.3 87.7 92.1 89.1 87.3 84.2 93.3 90.3 94.1 77.9 82.1 81.7 79.7 81.2 81.5 74.1 84.1 78.5

2 3 2 3 2 3 2 4 5 2 3 4 3 4 2 4 5 2

101.1 99.1 103.9 100.6 99.1 95.6 104.9 102.1 105.9 89.3 93.4 92.7 91.4 92.2 92.5 89.7 94.9 87.1

87.1 85.6 89.9 86.8 85.3 81.6 91.1 88.3 92.1 75.5 79.4 78.9 77.7 78.6 78.8 75.9 81.1 72.3

we also briefly examined the chemistry of benzylperoxy radical, because it is a hydrocarbon analog for methylated heteroaromatics and a more common experimental target. Calculations at the CBS-QB3 level closely replicated toluene’s experimentally determined geometry, spectroscopic information, and BDE; additionally, DFT (B3LYP) calculations replicated the qualitative trends predicted by the CBSQB3 calculations. Quantitatively, the DFT calculations consistently underpredicted the BDEs and reaction energies relative to CBS-QB3. The reactivity of methyl- and ethyl-substituted azabenzenes was explored by calculation of the homolytic BDEs and free energies for alkyl C—H hydrogen atom and alkyl side-chain loss. These values were analyzed as a function of heteroatom, ring size, side chain length, spin density, and temperature. Furthermore, the impact on the thermodynamic values derived from the harmonic oscillator approximation was analyzed, by treating side-chain torsions as hindered rotors. At 298 K, loss of hydrogen atom to form a benzylic-like radical is roughly 10 kcal/mol more favorable than loss of the alkyl group, due to the electron delocalization possible for the benzylic-like radical, regardless of ring size or heteroatom; this trend is consistently exhibited over a wide temperature range (T = 298–2000 K) although both reactions become increasingly favorable with increasing temperature, due to entropic effects on the free energy (Fig. 17). Spin densities for heteroaromatic radicals correlate well with the BDE values: the more diffuse the spin density, the lower the corresponding BDE. This is most dramatically evident for the five-membered ring heteroaromatics. Both the harmonic oscillator and hindered rotor treatments give comparable values for reaction enthalpies and free energies. The ethyl derivatives have lower reaction enthalpies and free energies than methyl derivatives.

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90

80

20

20

10

10

0

0

13

10

70

50

30

30

00 15 00 18 00 20 00

30

0

30

00

40

0

40

0

50

0 70 0 10 00 13 00 15 00 18 00 20 00

50

0

60

ΔG rxn (kcal/mol)

70

60

50

ΔG rxn (kcal/mol)

70

80

(b)

(a)

Temperature (K)

Fig. 17 Variation of reaction free energy (kcal/mol) with temperature (K) for alkyl C—H hydrogen atom loss in (a) five-membered ring methyl-substituted heteroaromatic rings and (b) six-membered ring methyl-substituted heteroaromatic rings.

With respect to predicting the chemistry of larger heteroaromatic systems (such as those in coal), these calculations suggest that both hydrogen atom and alkyl group loss can contribute to the combustion of coal in initiation reactions; the subsequent oxidation pathways of both the aromatic peroxy radicals (previously explored197,198) and the alkylated aromatic peroxy radicals are of interest. Reactivity will likely increase with increasing alkylation of a subunit, and the azabenzene units are more likely to react than the five-membered heteroaromatic rings. The initial steps of radical formation are expected to become more favorable at higher temperatures, primarily due to entropic considerations. Oxidative decomposition of alkylated heteroaromatics When alkylated heteroaromatic radicals are formed, these species can rapidly react with O2 to form various peroxy radicals. Again, it seems likely that these resultant reactive species will have an impact on overall combustion processes. Benzylperoxy radical was initially explored computationally to obtain a qualitative picture of peroxy radical decomposition for species containing both alkyl and aromatic components; these findings were then extended to peroxy radicals of methyl- and ethylsubstituted azabenzylperoxy radicals.205 The alkyl chains of these species are large enough than conformeric considerations are a concern. We performed calculations to analyze rotational profiles for alkyl group torsions to isolate the most stable conformation. In all cases, rotation of and alkyl dihedral angle occurred with a

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Fig. 18 Unimolecular pathways for 2-picolinylperoxy radical, a representative peroxy radical for alkylated azabenzenes.

barrier less than 5 kcal/mol. As shown in Fig. 18, alkylated heteroaromatic peroxy radicals (2) can revert to reactants (1) or cyclize via attack of the side chain’s terminal O at a position ipso (3) or ortho (4) to the alkylperoxy chain. Oxygen atom loss can occur (5). Given the presence of the exocyclic methylene group, isomerization via H-atom transfer (6) can occur from a methylene carbon to the terminal peroxy oxygen; this cannot occur in the nonalkylated parent compounds. Additionally, entire side chain loss (as dioxirane in the case of methyl-substituted aromatics) is possible (7). Calculated kinetic and thermodynamic values for unimolecular decomposition of benzylperoxy radical and 2-picolinylperoxy radical (2-methylpyridinylperoxy radical) at 298 K are presented in Table 4 (representative structures shown in Fig. 18). The energies for decomposition of these two compounds are similar: bicyclic ring formation at the carbon ortho to the alkyl group (pathway 4) occurs with a barrier of 30 kcal/mol; cyclization at the carbon ipso to the alkyl group (pathway 3) occurs with a barrier of 35–40 kcal/mol; internal H-atom transfer (pathway 6) occurs with a barrier of 40 kcal/ mol. The added functionality afforded by the nitrogen atom in 2-picolinylperoxy radical allows a cyclization resulting in a stable N—O radical after O—O bond scission (pathway 4b), which has potential for yielding precursors to NOx chemistry; however, this process has a substantially high barrier (50 kcal/mol) and is unlikely to be a factor at atmospheric temperatures. While the qualitative trends were replicated between benzylperoxy radical and 2-picolinylperoxy radical, the latter has slightly higher reaction barriers and energies in nearly every case, which can be attributed to the relative stabilization of the allylic systems in the cyclized derivatives. CBS-QB3 calculations helped demonstrate that the DFT (B3LYP) approach provides good qualitative predictions for the energetic trends. The trends for 2-picolinylperoxy radical were similar for other picolinylperoxy radicals, as well as peroxy radicals of alkylated diazabenzenes (i.e., pyridazine, pyrimidine, and pyrazine). It was observed that the presence of a second ring nitrogen has little effect on either the identity or energetics of the preferred pathways. Likewise, the ethyl

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Table 4 Comparison of reaction pathway energetics (kcal/mol) for benzylperoxy radical and 2-picolinylperoxy radical at 298 K, via B3LYP/6–311þG**//B3LYP/6–31G*. Numbers refer to pathways depicted in Fig. 18. All enthalpies and free energies are relative to the peroxy radical (2); when preceded by TS, the relative data are the enthalpy and free energies of activation Benzylperoxy radical B3LYP 

2!1 TS(2–3) 2!3 TS(2–4a) 2 ! 4a TS(2–6) 2!5 2!6 2!7

CBS-QB3 

DH

DG

–16 35.7 33.2 33.5 35.6 37.9 56.5 –33.9 53.7

–5.5 37.2 34.2 35.6 23.7 38.9 47.3 –35.4 42.9

DH



Experiment

DG



–22.7 30.1 25.8 29.7 14.6 38.4 61.9

–12.2 31.7 26.9 31.8 15.1 39.2 52.5

55.8

45.3

a

DH

DG

–20.1b, –21.8c

–12.2

a

2-Picolinylperoxy radical B3LYP

2!1 TS(2–3) 2!3 TS(2–4a) 2 ! 4a TS(2–4b) 2 ! 4b 2!5 TS(2–6) 2!6 2!7

CBS-QB3

DH

DG

DH

DG

16.3 37.1 34.2 30.4 20.6 47.1 –5.8 61.6 38.0 –24.7 49.2

6.5 39.4 35.7 33.2 23.2 49.8 –4.1 53.0 39.4 –33.6 38.7

22.7 31.0 26.1 25.6 13.2

13.2 33.4 27.9 28.7 16.0

37.6 –23.7

39.2 –32.4

Where applicable, 4a refers to cyclization at a ring carbon and 4b refers to cyclization at nitrogen. a Geometry could not be optimized. b Reference 161. c Reference 162.

analogs favored similar reaction pathways, with energies of activation and reaction varying by only 2 kcal/mol between the methyl and the ethyl analogs. When present, the disparities in the reaction energetics were rationalized via consideration of the inductive effects caused by the nitrogen atom(s) and varying amounts of geometric strain introduced in the rearrangement pathways. Fig. 19 shows that with increasing temperature, formation of 1, 5, and 7 benefit entropically and become substantially exoergic reactions. Formation of 3 and 4a exhibit little entropic benefit. As shown in Fig. 20, the reaction barriers similarly

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80 60

ΔGreaction (kcal/mol)

40 20 0 –20

0

1000

500

1500

2000

2500

×

–40

× ×

–60

×

–80

× ×

–100

×

–120

×

–140

T (K)

Fig. 19 DGrxn versus temperature for the unimolecular pathways of 2-picolinylperoxy radical. 2 ! 3 denoted by open diamond; 2 ! 4a denoted by solid square; 2 ! 4b denoted by open triangle; 2 ! 5 denoted by dash; 2 ! 6 denoted by symbol ; 2 ! 7 denoted by solid diamond. All energies calculated at the B3LYP/6–311þG(d,p)//B3LYP/6–31G(d) level of theory.

50 45 40 ΔGactivation (kcal/mol)

× ×

35

×

30

×

25

× ×

20

×

15

×

10 5 0

500

1000

1500

2000

2500

T (K)

Fig. 20 DGactivation versus temperature for unimolecular pathways for 2-picolinylperoxy radical. Activation energy for 2 ! 3 denoted by open diamond; activation energy for 2 ! 4a denoted by solid square; activation energy for 2 ! 4b denoted by open triangle; activation energy for 2 ! 6 denoted by symbol  . All energies calculated at the B3LYP/ 6–311þG(d,p)//B3LYP/6–31G(d) level of theory.

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demonstrate minimal entropic effects. Formation of 4b has a similarly small entropic benefit but maintains a high reaction barrier, such that any NOx formation via this process is expected to be minimal except at extremely high temperatures. Formation of 6 exhibits a precipitous drop in its barrier and is the dominant processes at T  1250 K. This last reaction is of considerable interest given its relevance to both lowand high-temperature combustion, since it essentially constitutes a dissociative rearrangement pathway. Both the direct and the indirect decomposition intermediates (1, 5, 6, and 7) will play a substantial role in combustion as temperatures rise above 298 K, as peroxy radicals, themselves ROS, demonstrate the potential to generate reactive O-atom and hydroxyl radical. The overall chemistry of alkylated azabenzylperoxy radicals was consistent regardless of alkyl substitution or number of nitrogen atoms. The picolinylperoxy radicals provide excellent models for the chemistry exhibited by this larger class of species. Moreover, aromatic hydrocarbons can themselves predict several aspects of this chemistry, the exception being those processes involving oxidation or rearrangement with ring nitrogens. When this approach is extended to alkylated derivatives of the five-membered heteroaromatic rings,206 the energetic trends can be considered in light of ring size and heteroatom effects. In general, the same six pathways are available, such that rearrangements and dissociations are observed to contribute to the chemistry of these species. Many trends are consistent between both sets of alkylated heteroaromatics: increasing alkyl substitution leads to small stabilization of reaction barriers and energies. Increasing temperature leads to a shift in preference for the dissociation pathways. However, a difference was observed for smaller rings: while cyclizations are favored pathways at 298 K regardless of parent ring size, alkyl-substituted heteroaromatic peroxy radicals of the five-membered rings cyclize in such a manner as to generate an allylic radical system (cyclizing at a position either ipso or ortho to the side chain). On the other hand, alkyl-heteroaromatic peroxy radicals of the sixmembered rings preferred to form five-membered rings via ortho cyclization, due to their increased stability over spirodioxetane structures (which would result from ipso cyclization). This general trend is represented in Fig. 21.

Fig. 21 Qualitative depiction of favorable cyclization pathways for representative peroxy radicals of methyl heteroaromatics (top, pyridine and bottom, furan). Cyclization for the alkylated six-membered heteroaromatics is driven by the thermodynamic stability of the resulting ring, while cyclization for the alkylated five-membered heteroaromatics is dictated by which pathway allows the generation of a stable allylic radical system.

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Heteroatom identity also has an impact on the reactions of interest. Oxygen, nitrogen, and sulfur atoms affected heteroaromatic chemistry to slightly different quantitative extents (though qualitative trends were consistent), and the effect of multiple heteroatoms (as in oxazole) led to small differences in reaction energetics. Notably, the pyrrole, oxazole, and thiophene analogs reacted to form various species wherein their heteroatoms are oxidized, and these effects were unique to each heterocycle. Thus, data for aromatic hydrocarbon species are useful in estimating the energetics of these species but cannot completely predict their chemistry, due to the propensity of increasing functionalization to lead to different products and pathways. Relative to nonalkylated heteroaromatic peroxy radicals,197,198 we observed that alkylated species demonstrated a greater affinity for intramolecular reactions, due to the length and flexibility of their side chains. Dissociative reactions were less favorable for alkylated derivatives, due to a reduction of aromatic character. Our studies have provided important details regarding the oxidative decomposition of alkyl-substituted heteroaromatic rings: in particular, pathways originating with peroxy radicals derived from the heteroaromatic rings and temperature effects on these pathways. We have shown that the alkylated five-membered heteroaromatics demonstrate certain unique tendencies in their reactivity. Given the differences in reactivity between the five- and the six-membered heteroaromatics, as well as between the alkylated and the nonalkylated heteroaromatics, it seems likely that the chemical behavior of coal could vary somewhat depending on the abundance and nature of the cyclic subunits present in its structure. However, these discrepancies are mainly limited to the variety of products formed via combustion rather than the overall kinetics and thermodynamics of the relevant processes.

4

Future challenges in combustion chemistry

In closing this review, it seems logical to highlight some of the most recent progress in combustion chemistry. Alternate forms of energy and methods of combustion are continually being developed, thereby continuously invoking new challenges for experimentalists and theoreticians. Some alternatives involve new combustion processes, as in the case of homogenous-charge compression-ignition (HCCI)207 chemistry, which depends on autoignition (low temperature) chemistry within a homogenous gas mixture (unlike a SI engine). This new combustion method has the potential to increase fuel efficiency and decrease harmful emissions.208,209 One significant issue with HCCI is adequately controlling fuel ignition during the compression stroke of the engine; this is essentially a kinetics problem, requiring understanding of the reaction rates and mechanisms of the radical species generated by the fuel. Thus, reactions normally associated with low-temperature combustion and often treated cursorily in oxidative mechanisms can have a greater impact on combustibility. In this review, we are most interested in the implications of the changing energy landscape for the reactive pathways involved. As the nature of the fuels used in everyday life change, the reactive radical intermediates formed in their combustion will also change. We will examine the context of two classes of relevant fuel developments.

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FUEL ADDITIVES

Oxygenated molecules have long been used to augment gasoline formulations. For instance, since the elimination of tetraethyl lead (Pb(C2H5)4) as an additive. oxygenates have been used as gasoline additives to reduce harmful CO emissions and increase a fuel’s octane rating. The most prominent oxygenate has been methyl tert-butyl ether (MTBE); however, this additive has since been flagged as a potential carcinogen and odorous component that can taint water supplies.210 Moreover, MTBE necessitates the use of the industrial byproduct, isobutene, for its generation, further competing with crude oil supplies for a value-added product. As the downsides to MTBE use have become more apparent, ethanol has emerged as an attractive alternative oxygenate. Ethanol does not demonstrate the negative health effects of MTBE, and moreover, it can be produced from renewable (biomass) materials.211 Compared to MTBE, ethanol has a higher oxygen-to-carbon ratio. Ethanol has been used primarily as a gasoline additive; in fact in the years since the Energy Policy Act of 1992,212 a shift has occurred, such that current manufacturers are building cars with the capacity to run on fuel blends of 85 and 95% ethanol, along with a small amount of gasoline – these fuels are often referred to as E85 or E95, respectively. (However, some engine modifications are necessary for vehicles to run effectively on blended fuels with greater than 20% ethanol.) This advance has been hampered by the general lack of availability of ethanol at fueling stations, a shortage that is being gradually remedied. Currently, the cost benefits of using ethanol over gasoline are substantial, but the availability of the former is still limited.213 Ethanol itself demonstrates several drawbacks as an alternative fuel, despite its increasing availability. Most notably, ethanol absorbs water. Thus, it cannot travel through existing gasoline pipelines, as the water could subsequently separate and freeze during colder temperatures, possibly bursting the pipelines; moreover, ethanol is also corrosive.214 Thus, ethanol has to be transported via other means (a fact which, ironically, generally necessitates the use of gasoline, diesel, or other fuel and results in a higher cost that negates one monetary benefit of using ethanol as a fuel). Moreover, ethanol evaporates relatively quickly, and so in warmer temperatures, it must be blended carefully.215 Butanol possesses the chemical benefits of ethanol while avoiding its drawbacks. Its use is considerably less temperature-sensitive: it is six times less evaporative than ethanol, and it is not corrosive, so can be shipped via existing fuel pipelines.216 In terms of fuel benefits, due to its higher number of carbons, it has higher energy content (110,000 Btu/gallon) than ethanol (84,000 Btu/gallon) and is comparable to gasoline (115,000 Btu/gallon); it can also be generated more easily from biomass than can ethanol.217 Unlike ethanol, it can be used as a direct replacement for gasoline rather than as an additive therein, since butanol’s air/fuel ratio is comparable to that of gasoline.218 Finally, butanol’s primary combustion byproduct is carbon dioxide; it avoids formation of the pollutants NOx, SOx, and CO.219 Thus, butanol is a potential fuel of substantial interest. Studies have been completed specifically on combustion processes of ethanol and butanol, and several of the peroxy and oxy radical species have been examined.95–108

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Marinov proposed an exhaustive mechanism for ethanol’s combustion.220 Cavalli et al. examined the initial reaction of HO• radical with 1-butanol by FTIR spectroscopy,221 and Chen et al. studied the anodic oxidation of this species.222 Recent experiments by McEnally and Pfefferle led them to propose that butanol combustion primarily occurs via a complex fission reaction rather than H-atom abstraction.223 While we have discussed the historical background of the main gasoline oxygenates, it is worthwhile to note that several other species have been discussed in this context as well. Notably, it has been shown that simply adding oxygen to a given combustion environment does not in itself achieve soot reduction; the role of the oxygen within the structure of the oxygenate plays a major part. For instance, Westbrook et al. completed a modeling study on the effect of various oxygenated hydrocarbons on soot production.224 They saw a significant reduction in the number of alkynyl radicals serving as soot precursors in a diesel flame via the inclusion of alcohol (ROH) and ether (ROR0 ) additives, but noted a lesser effect when esters (RCO2R0 ) were included. This is due to the fact that esters can readily fragment to generate CO2, so that their oxygen atoms are not involved in the processes that can affect soot reduction. Similarly, Sinha and Thomson225 studied three C3 oxygenated hydrocarbons – isopropyl alcohol, dimethoxy methane (DMM), and dimethyl carbonate (DMC) – all in comparison to propane, noting different effects for each. DMM and DMC lack C—C bonds. Thus, the concentration of the alkenyl and alkynyl radicals necessary to form benzene and larger aromatic hydrocarbons is correspondingly smaller in these species and soot production drops. Nag et al.226 have postulated that the identity of a given oxygenate contributes to the resultant balance of CO and CO2 in the emission pool; CO has a greater propensity for PAH reduction than does CO2. The implications for power output and efficiency are important as well, since a significant amount (33%) of the eventual heat is generated by the conversion of CO to CO2 in the combustion environment (usually mediated by HO• radical). The structure of an oxygenate has notable implications for both its reactivity and its propensity for the formation of emissions.

BIODIESEL

Having alluded briefly to biomass compounds in the previous section, as sources for ethanol and butanol, we turn now to the potential of these species as fundamental alternatives to fossil fuels. In 1912, Rudolf Diesel presciently stated,227 ‘‘The use of vegetable oils for engine fuels may seem insignificant today. But such oils may in the course of time become as important as petroleum and the coal tar products of the present time.’’ Biodiesel can be produced from a variety of sources,228 including animal fats (on small scales), algae, and vegetable oils: the fuel itself is produced via esterification of these lipids with methanol or ethanol. In terms of the overall benefits and sources of fuels, several reviews are currently available. As early as 1987, Schwab et al. reported on the potential of vegetable oils for forming diesel fuels.229 More recently, Graboski and McCormick,230 Srivastava and Prasad,231 and Lin et al.232 have reviewed the fuel properties (emissions, engine performance, etc.) of

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these species. Ma and Hanna have compiled a review of the general benefits of biodiesel chemistry.233 The emergence of biofuels as potential energy sources demands a mechanistic approach comparable to those currently in use for hydrocarbon compounds. Biofuels are commonly esters (RCO2R0 ) of large long-chain fatty acids, so their reactivity will depend on more factors than these extant mechanisms. Kulkarni and Dalai have summarized mechanistic aspects of biodiesel chemistry, with respect to waste cooking oil.234 Thermolytic decomposition can yield alkanes, alkenes, ketones, esters, and small acids. Oxidative decomposition leads to highly functionalized peroxy radicals, which then can add a hydrogen atom to form hydroperoxides, ultimately decomposing to aldehydes, hydrocarbons, and acids; additionally, these peroxy radicals might dimerize or oligomerize if excess oxygen is present. Hydrolytic reactions comprise an additional possibility: triglycerides readily decompose to glycerol, monoglycerides, diglycerides, and free fatty acids (FFA), in the presence of water. Zhenyi et al. have completed thermodynamic calculations on the pyrolysis of vegetable oils, postulating that, for a given ester, the key initiation step proceeds via breaking of the alkyl (sp3) C—O bond.235 We can summarize the different possible initiation routes for a given triglyceride, familiar from our previous discussions (Fig. 22). Even from this basic view, it is clear that several classes of reactive radical intermediates will play a role in biofuel combustion: alkyl (sp3) and alkenyl (sp2) radicals are readily formed, as well as peroxy radicals and functionalized derivatives of each of these original species.

Fig. 22 Initial reaction steps for fatty acid ester decomposition, represented by ethyl butanoate, which can decompose via thermolysis (pyrolysis), oxidation, or hydrolysis.

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Mechanism development traditionally depends on the inclusion of data for relevant, smaller compounds; that philosophy will apply in this case, where the reactions of both small hydrocarbons and small esters are of interest. Mechanisms for methyl butanoate (CH3CH2CH2CO2CH3) have been developed by Gail et al.236 and Curran et al.;237 the latter notes the NTC range demonstrated by this species, suggesting that future work may benefit by exploring comparisons to hydrocarbon chemistry. Good and Francisco have modeled the tropospheric oxidation mechanism of methyl formate (HCO2CH3).238 Metcalfe et al. explored unimolecular decomposition pathways and derived combustion mechanisms for methyl butanoate and ethyl propanoate, using shock-tube experiments and CBS-QB3 calculations.239 Recent studies by Glaude et al.,240 Schwartz et al.,241 and Sarathy et al.242 have provided experimental information on larger esters, in an effort to model further reactions with implications for biofuel combustion. This area is of significant interest, and it is expected that the next few years will provide a more thorough understanding of the important species and relevant mechanisms involved in biofuel combustion.

5

Conclusions

Throughout this review, we have provided a historical context for understanding combustion chemistry as it applies to some of the most fundamental hydrocarbon compounds. We have explored the implications of these processes for larger and increasingly functionalized molecules. These species can be used to model the chemistry of petroleum and coal, in addition to smaller hydrocarbon fuels. Combustion is a complex topic, such that both theoretical and experimental methods are useful in exploring chemistry with implications for high-temperature oxidation and lowtemperature atmospheric reactions. In particular, the master equation methods developed over the past few decades can consider the collective chemistry of thousands of elementary steps to predict the overall oxidation of a given fuel; the kinetics and thermodynamics of each of these elementary steps can be generated via experiment or computational modeling. ROS and other radical intermediates dictate the oxidative decomposition of fuels. We have noted that peroxy radical intermediates provide an enormous amount of flexibility in the combustion of a given compound, specifically in the unimolecular steps available to that compound. In an instructive display of the interaction of experimental and theoretical techniques, rearrangement pathways of the peroxy radicals have been modeled computationally and provide justification for several unexpected products. At the start of the twenty-first century, efforts are underway to decrease society’s dependence on fossil fuels. It is clear that alternate energy forms will bring with them their own sets of reactive radical intermediates and revisit the important intermediates seen from smaller model compounds, as we consider future challenges in combustion chemistry. We expect that advances in experimental techniques and computational approaches will correspondingly be developed in the years ahead.

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Acknowledgments We are grateful to Dr. Timothy Barckholtz (ExxonMobil Research and Engineering) for helpful discussions and for providing Fig. 3. We extend sincere thanks to Dr. Donald Burgess, Jr. (NIST) for his help in the editing of this review.

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