Alkyl hydroperoxides determination

Disproportionation reaction 7 might be expected to be thermoneutral in the gas phase and perhaps less so in the liquid phase where there is the possibility of hydrogen-bonding. Only for gas phase dimethyl peroxide is the prediction true, where the reaction enthalpy is —0.2 kJmoD. The liquid phase enthalpy of reaction is the incredible —61.5 kJmoD. Of course, we have expressed some doubt about the accuracy of the enthalpy of formation of methyl hydroperoxide. For teri-butyl cumyl peroxide, the prediction for thermoneutrality is in error by about 6 kJmor in the gas phase and by ca 9 kJmoD for the liquid. The enthalpy of reaction deviation from prediction increases slightly for tert-butyl peroxide — 14kJmol for the gas phase, which is virtually the same result as in the liquid phase, — 19kJmol . The reaction enthalpy is calculated to be far from neutrality for 2-fert-butylperoxy-2-methylhex-5-en-3-yne. The enthalpies of reaction are —86.1 kJmoD (g) and —91.5 kJmol (Iq). This same species showed discrepant behavior for reaction 6. Nevertheless, still assuming thermoneutrality for conversion of diethyl peroxide to ethyl hydroperoxide in reaction 7, the derived enthalpies of formation for ethyl hydroperoxide are —206 kJmoD (Iq) and —164 kJmoD (g). The liquid phase estimated value for ethyl hydroperoxide is much more reasonable than the experimentally determined value and is consistent with the other n-alkyl hydroperoxide values, either derived or accurately determined experimentally. 

As was the case for the alkyl hydroperoxides in reaction 4, the enthalpies of the oxy-gen/hydrocarbon double exchange reaction 8 for dialkyl peroxides are different depending on the classification of the carbon bonded to oxygen. For R = Me, Et and f-Bu, the liquid phase values are —4, 24.6 and 52.7 kJmoR, respectively, and the gas phase values are 0.1, 25.7 and 56.5 kJmoR, respectively. For the formal deoxygenation reaction 9, the enthalpies of reaction are virtually the same for dimethyl and diethyl peroxide in the gas phase, —58.5 0.6 kJ moR. This value is the same as the enthalpy of reaction of diethyl peroxide in the liquid phase, —56.0 kJ moR (there is no directly determined liquid phase enthalpy of formation of dimethyl ether). Because of steric strain in the di-ferf-butyl ether, the enthalpy of reaction is much less negative, but still exothermic, —17.7 kJmol (Iq) and —19.6 kJmol (g). 

Alkyl halides, hydroperoxide synthesis, 327-8 Alkyl hydroperoxides anion ligands, 114-19 covalent radii, 114, 118-19 dihedral angles, 119 geometric parameters, 115-8 tetrahedral distortion, 119 artemisinin formation, 133-4 chlorotriorganosilane reactions, 779-83 crystal structure, 105-14 anomeric effect, 110-11 geometric parameters, 106-9 hydrogen bonding, 103-5, 111-14 tetrahedral distortion, 110 determination, 674... 

Cobalt(ll)-EDTA complex, hydrogen peroxide determination, 628, 639 Cobalt(ll)-hexacyanoferrate, hydrogen peroxide determination, 651 Cobalt(lll)-phthalocyaninetetrasulfonate, hydroperoxide determination, 677 CocrystaUization, alkyl hydroperoxides-ether, 111, 113... 

Competition between metal ion-induced and radical-induced decompositions of alkyl hydroperoxides is affected by several factors. First, the competition is influenced by the relative concentrations of the metal complex and the hydroperoxide. At high concentrations of the hydroperoxide relative to the metal complex, alkoxy radicals will compete effectively with the metal complex for the hydroperoxide. Competition is also influenced by the nature of the solvent (see above). Contribution from the metal-induced reaction is expected to predominate at low hydroperoxide concentrations and in reactive solvents. The contribution from the metal-induced decomposition to the overall reaction is readily determined by carrying out the reaction in the presence of free radical inhibitors, such as phenols, that trap the alkoxy radicals and, hence, prevent radical-induced decomposition.129,1303 Thus, Kamiya etal.129 showed that the initial rate of the cobalt-catalyzed decomposition of tetralin hydroperoxide, when corrected for the contribution from radical-induced decomposition by the... 

Determination of several atmospheric oxidant species is critical to understanding gaseous and aqueous processes leading to acidic deposition. Hydrogen peroxide has a high Henry s Law solubility and must be measured in gaseous and aqueous atmospheric samples to better understand wet deposition processes. In contrast, measurements of ozone and peroxyacyl nitrates (PANs) (and probably alkyl hydroperoxides and peracids) usually need to be made only in the gas phase due to their low aqueous solubility (14). 

Sulfides react faster with hydrogen peroxide and alkyl hydroperoxides than do alkenes. For this reason, transition metal catalysts are rarely necessary, but these reactions are acid catalyzed and first order in both sulfide and peroxide. The acid (HX) can be as weak as alcohol or water but the "effectiveness (of the oxidation) is determined by the pXa of the acid. Sulfides also react faster with peroxides than do ketones (see the Baeyer-Villiger reaction, sec. 3.6). Formation of the sulfone in these reactions is straightforward, but requires more vigorous reaction conditions. It is usually easy to isolate the sulfoxide from oxidation of a sulfide. Direct conversion of a sulfide to a sulfone requires excess peroxide and vigorous reaction conditions (heating, long reaction times, more concentrated peroxide). 

Total concentrations of detectable stable products (cyclohexanol and cyclohexanone) in the chromatogram of the sample unreduced with PPhj will be (C. + ACoiT" and (Cam ACo, ) " , respectively. Triphenylphosphine reduces the alkyl hydroperoxide to yield, quantitatively, cyclohexanol. So the concentrations of ol and one determined by GC in the sample treated with PPhj will be (C , + Cpnd) and (Co e)" , respectively. The real amount of ketone present in the reaction mixture may thus be determined by measuring the concentration (Cone) " of this product in the reduced sample. 

However, if there is only one channel of the ol and one formation in the course of the oxidation, i.e., gradual decomposition of the alkyl hydroperoxide to produce both products again in approximately equal amounts, the concentrations of ol and one in the chromatogram of the untreated sample should be equal. This situation has been noticed in many cases of alkane oxidation. If the concentrations of ol and one, determined by GC, in the untreated sample are different, it testifies that the real amounts of products are not equal and either the alkyl hydroperoxide decomposes to produce predominantly the alcohol (or, on the contrary, the ketone) or there is an additional channel leading to cyclohexanol or cyclohexanone. This third simplified (without determination of ) method can give relatively precise values of the real concentrations of the alkyl hydroperoxide (as well as cyclohexanol and cyclohexanone) only if all conditions mentioned above are valid. For example, if the chromatogram of a solution before the reduction exhibits two peaks of approximately equal area for cyclohexanol and cyclohexanone, and after the reduction only cyclohexanol is determined by the GC, these data testify that only cyclohexyl hydroperoxide is present in the solution and its concentration can be determined precisely. 

It should be noted, that the approximate equality of alcohol and ketone yields determined in the untreated sample generally by no means can be interpreted as evidence for the presence of the alkyl hydroperoxide in the reaction... 

Overall the vinyl, phenyl, and ethynyl-alkyl hydroperoxides and methylperoxides do not exist as stable species under atmospheric conditions and its life time scales of seconds. Determination of the enthalpies of formation provides analysis of the relative bond strengths in these species and subsequent information concerning the reactions of peroxide radicals, which are important in both combustion and atmospheric chemistry of hydrocarbons. Reactions of... 

The pH dependencies of kiy for the bimolecular reactions of (l)MnlH(X)2 with three alkyl hydroperoxides are shown in Figure 3. For (l)MnlH(X)2, the kinetics of the reactions display pH dependence consistent with a scheme involving the critical formation of two hydroperoxide-coordinated manganese(III) intermediates at intermediate and high pH values. At low pH the the reaction of (l)MnlH(X)2 with hydroperoxides could not be determined over spontaneous decomposition of hydroperoxide. The kinetic expression was derived from the pathway shown in eq 16 with a steady state assumption in intermediates IIH and II". 

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