Ethane reactions, calculations

The quantum-chemical calculation of the activation energies of reactions EtOO + EtH and EtOO + EtOH by the DFT method supports this result [54]. For the reaction of EtOO with ethane, the calculated E value is 82 kJ mol-1, AH = 76 kJ mol-1, and Ee0 = 56 kJ mol-1. For the reaction of EtOO with ethanol is = 52.7. The structure of TS is presented in Figure 7.1. 

In order to explain the data of Aronowitz et al (12) and previous shock—tube and flame data, Westbrook and Dryer (12) proposed a detailed kinetic mechanism involving 26 chemical species and 84 elementary reactions. Calculations using tnis mechanism were able to accurately reproduce experimental results over a temperature range of 1000—2180 K, for fuel—air equivalence ratios between 0.05 and 3.0 and for pressures between 1 and 5 atmospheres. We have adapted this model to conditions in supercritical water and have used only the first 56 reversible reactions, omitting methyl radical recombinations and subsequent ethane oxidation reactions. These reactions were omitted since reactants in our system are extremely dilute and therefore methyl radical recombination rates, dependent on the methyl radical concentration squared, would be very low. This omission was justified for our model by computing concentrations of all species in the reaction system with the full model and computing all reaction rates. In addition, no ethane was detected in our reaction system and hence its inclusion in the reaction scheme is not warranted. We have made four major modifications to the rate constants for the elementary reactions as reported by Westbrook and Dryer (19) ... 

The qualitative interpretation of the ISM method can be illustrated by Figure 11.18. Figure 11.18a shows the equi-energy curves that pertain to the transfer of H between methyl and methane and ethyl and ethane. The curves nearly overlap, but the slightly weaker C-H bond in ethane is characterized by a smaller force constant and leads to a somewhat smaller barrier. The calculated barriers are 14.6 and 14.3 kcal/mol, respectively. The methyl-ethane reaction, shown in Figure 11.18b, is exothermic and there is a much more substantial shift in the curves. The calculated barrier is 12.4 kcal, compared with the experimental value of 11.5. Thus, the calculation moves the barriers in the right direction, although it does not reproduce the entire effect that is observed experimentally. 

Irradiation of ethyleneimine (341,342) with light of short wavelength ia the gas phase has been carried out direcdy and with sensitization (343—349). Photolysis products found were hydrogen, nitrogen, ethylene, ammonium, saturated hydrocarbons (methane, ethane, propane, / -butane), and the dimer of the ethyleneimino radical. The nature and the amount of the reaction products is highly dependent on the conditions used. For example, the photoproducts identified ia a fast flow photoreactor iacluded hydrocyanic acid and acetonitrile (345), ia addition to those found ia a steady state system. The reaction of hydrogen radicals with ethyleneimine results ia the formation of hydrocyanic acid ia addition to methane (350). Important processes ia the photolysis of ethyleneimine are nitrene extmsion and homolysis of the N—H bond, as suggested and simulated by ab initio SCF calculations (351). The occurrence of ethyleneimine as an iatermediate ia the photolytic formation of hydrocyanic acid from acetylene and ammonia ia the atmosphere of the planet Jupiter has been postulated (352), but is disputed (353). 

An example of the application of molecular mechanics in the investigation of chemical reactions is a study of the correlation between steric strain in a molecule and the ease of rupture of carbon-carbon bonds. For a series of hexasubstituted ethanes, it was found that there is a good correlation between the strain calculated by the molecular mechanics method and the rate of thermolysis. Some of the data are shown in Table 3.3. 

In order to generate the starting material for a polymer that is used in water bottles, hydrogen is removed from the ethane in natural gas to produce ethene in the catalyzed reaction C,H6(g) H,(g) + C,ll4(g). Use the information in Appendix 2A to calculate the equilibrium constant for the reaction at 298 K. (a) If the reaction is begun by adding the catalyst to a flask containing C,H6 at 10.0 bar, what will be the partial pressure of the C,H4 at equilibrium (b) Identify three steps the manufacturer can take to increase the yield of product,... 

Transient computations of methane, ethane, and propane gas-jet diffusion flames in Ig and Oy have been performed using the numerical code developed by Katta [30,46], with a detailed reaction mechanism [47,48] (33 species and 112 elementary steps) for these fuels and a simple radiation heat-loss model [49], for the high fuel-flow condition. The results for methane and ethane can be obtained from earlier studies [44,45]. For propane. Figure 8.1.5 shows the calculated flame structure in Ig and Og. The variables on the right half include, velocity vectors (v), isotherms (T), total heat-release rate ( j), and the local equivalence ratio (( locai) while on the left half the total molar flux vectors of atomic hydrogen (M ), oxygen mole fraction oxygen consumption rate... 

While alkane metathesis is noteworthy, it affords lower homologues and especially methane, which cannot be used easily as a building block for basic chemicals. The reverse reaction, however, which would incorporate methane, would be much more valuable. Nonetheless, the free energy of this reaction is positive, and it is 8.2 kj/mol at 150 °C, which corresponds to an equihbrium conversion of 13%. On the other hand, thermodynamic calculation predicts that the conversion can be increased to 98% for a methane/propane ratio of 1250. The temperature and the contact time are also important parameters (kinetic), and optimal experimental conditions for a reaction carried in a continuous flow tubiflar reactor are as follows 300 mg of [(= SiO)2Ta - H], 1250/1 methane/propane mixture. Flow =1.5 mL/min, P = 50 bars and T = 250 °C [105]. After 1000 min, the steady state is reached, and 1.88 moles of ethane are produced per mole of propane consmned, which corresponds to a selectivity of 96% selectivity in the cross-metathesis reaction (Fig. 4). The overall reaction provides a route to the direct transformation of methane into more valuable hydrocarbon materials. 

Carbonyl group of the aldehyde decreases the BDE of the adjacent C—H bond. This is due to the stabilization of the formed acyl radical, resulting from the interaction of the formed free valence with Tr-electrons of the carbonyl group. For example, DC—H = 422kJmol 1 in ethane and D( n 373.8 kJ mol 1 in acetaldehyde. The values of Dc H in aldehydes of different structures are presented in Table 8.1. In addition, the values of the enthalpies of acylperoxyl radical reactions with aldehydes were calculated (D0 H= 387.1 kJ mol-1 in RC(0)00 H). 

Lehnert and Tuczek further studied end-on terminal coordination by density functional theory (DFT) calculations on the compounds [Mo(N2)2(dppe)2], [MoF(NNH)(dppe)2], and [MoF(NNH2)(dppe)2]+, where dppe= 1,2-bis(diphenyl-phosphino)ethane.50 They proposed a reaction scheme, shown in reaction 6.13, for asymmetric dinitrogen reduction and protonation. The end-on model favored by Lehnert in reference 50, as shown in reaction 6.13, appears to be a less thermodynamically unfavorable pathway, at least to reach the M-NNH3 intermediate. Step 1 produces a metal-attached diazenido ion (NNH-), step 2 produces a hydrazido ion (NNH2 ), and step 3 produces a hydrazidium ion (NNHj). 

Consider the gas-phase decomposition of ethane (A) to ethylene at 750°C and 101 kPa (assume both constant) in a PFR. If the reaction is first-order with kA = 0.534s-1(Fro-ment and Bischoff, 1990, p. 351), and r is 1 s, calculate /a- For comparison, repeat the calculation on the assumption that density is constant. (In both cases, assume the reaction is irreversible.)... 

At 252 °C based on kg/ks = 0.15 reaction (9) accounts for only 34 % of the ethane and 11 % of the ethylene. Reactions (6) and (7) are required to explain the concordance of results based on gas analysis and with those based on tetramethyl lead analysis. All observed orders and activation energies are consistent with this mechanism. If reaction (1) is the rate-controlling step in the initiation, the rate of this reaction can be calculated from... 

In some cases, the value given in the table depends on that calculated previously for some other bond. For example, to obtain ec-c, we combine the enthalpy of combustion of ethane, —1,588kJ mol , with the proper multiples of the AHm s in Equations (4.38)-(4.41) to obtain the enthalpy change for the reaction... 

The protodelithiation enthalpy of n-propyl lithium is very nearly the same as for the n-butyl species, —219 2 kJmoP. From reaction 10 with w-butyl lithium as the benchmark species and the enthalpies of formation of the hydrocarbons in their gaseous reference states, the enthalpy of formation of n-propyl lithium is calculated as ca —91 klmoP, a value consistent with that of —86 kJ moP derived from w-PrMgBr in an earlier section. If the reference state of n-butane is taken as the liquid instead, the enthalpy of formation of n-propyl lithium is ca —70 kJ moP, a value consistent with another previous derivation of ca —73 kJmoP. At least with respect to consistency with the enthalpies of formation in Table 1, the best reference state for ethane and propene is the gas it is not yet clear which is better for butane. 

In this paper selectivity in partial oxidation reactions is related to the manner in which hydrocarbon intermediates (R) are bound to surface metal centers on oxides. When the bonding is through oxygen atoms (M-O-R) selective oxidation products are favored, and when the bonding is directly between metal and hydrocarbon (M-R), total oxidation is preferred. Results are presented for two redox systems ethane oxidation on supported vanadium oxide and propylene oxidation on supported molybdenum oxide. The catalysts and adsorbates are stuped by laser Raman spectroscopy, reaction kinetics, and temperature-programmed reaction. Thermochemical calculations confirm that the M-R intermediates are more stable than the M-O-R intermediates. The longer surface residence time of the M-R complexes, coupled to their lack of ready decomposition pathways, is responsible for their total oxidation. 

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