Dimethyl ether (DME) has been proposed to be a promising alternative to conventional diesel fuel because of its favorable compression ignition property (high cetane number) and its soot-free combustion. A radical chain mechanism for hydrocarbon autoignition has been proposed for DME at low temperatures. In this mechanism, the chain initiation step consists of DME undergoing hydrogen abstraction by a highly reactive species (typically ·OH). The CH 3OĊH2 created in the initiation step then combines with O2; the subsequent CH3OCH2OO· radical is involved in a Lindemann-type mechanism, which can lead to the production of formaldehyde (CH2 = O) and ·OH. This concludes the chain-propagating step: the one ·OH produced then sustains the chain-reaction by creating another CH3OĊH2. A relatively stable intermediate (·CH2OCH2OOH), formed via isomerization of CH3OCH2OO· in the chain-propagation step, can combine with a second O2 to produce a radical (·OOCH2OCH2OOH) that can potentially decompose into two ·OH radical (and other products). This path leads to chain-branching and an exponential increase in the rate of DME oxidation. We have used spin-polarized density functional theory with the Becke-3-parameter Lee-Parr-Yang exchange-correlation functional to calculate the structures and energies of key reactants, intermediates, and products involved in (and competing with) the chain-propagating and chain-branching steps of low-temperature DME oxidation. In this article, Part I, we consider only the chain-propagation mechanism and its competing mechanisms for DME combustion. Here, we show that only certain conformers can undergo the isomerization to ·CH2OCH2OOH. A new transition state has been discovered for the disproportionation reaction ·CH2OCH 2OOH → 2CH2O + ·OH in the chain-propagating step of DME autoignition that is much lower than previous barriers. The key to making this decomposition pathway facile is initial cleavage of the O-O rather than the C-O bond. This renders all transition states along the chain-propagation potential energy surface below the CH3OĊH 2 + O2 reactants. In contrast with the more well-studied CH3ĊH2 (ethyl radical) + O2 system, the H-transfer isomerization of CH3OCH2OO· to ·CH2OCH2OOH in low-temperature DME oxidation has a much lower activation energy. This is most likely due to the larger ring strain of the analogous transition state in ethane oxidation, which is a five-membered ring opposed to a six-membered ring in dimethyl ether oxidation. Thus low-temperature ethane oxidation is much less likely to form the ·ROOH (where R is a generic group) radicals necessary for chain-branching, which leads to autoignition. Three competing reactions are considered: CH 3OĊ·H2 → CH2O + ·CH3; ·CH2OCH2OOH → 1,3-dioxetane + ·OH; and ·CH2OCH2OOH → ethylene oxide + HOO·. The reaction barriers of all these competing paths are much higher in energy (7-10 kcal/mol) than the reactants CH 3OĊH2 + O2 and, therefore, are unlikely low-temperature paths. Interestingly, an analysis of the highest occupied molecular orbital along the CH3OĊH2 decomposition path shows that electronically excited (1A2 or 3A2) CH2O can form; this can also be shown for ·CH2OCH2OOH, which forms two formaldehyde molecules. This may explain the luminosity of DME's low-temperature flames.
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