Oxidation-reduction rate law

Oxidant Reductant Rate law Medium Rate constant (25 °C) AH AS Ref... 

The rate of oxidation of isopropyl alcohol is also enhanced by the presence of glycolic acid (GA) as a co-oxidant. The rate law is complex but a kinetic term, first-order in [HCrOi"], [GA], and [ROH], indicates that the reduction of chromium(vi) through a termolecular complex must be more feasible than through either Cr -glycolic acid or Cr -alcohol complexes themselves. The reason for this preference is considered to be the avoidance of unstable intermediates, particularly Cr. This is accomplished by a one-step three-electron reduction of chromium(vi) to chroraium-(in) (Scheme 2). Two of the electrons are provided by the alcohol and the third by... 

By way of contrast, oxidation of the organic ligand in oxalatopentaamminecobalt-and p-aldehydobenzoatopentaamminecobalt(lll) is accompanied by reduction of the cobalt(IIl) centre in the case of one-equivalent oxidants, e.g. Ce(lV), but not in the case of two-equivalent oxidants e.g. CI2). The rate law is simple... 

COVALENT COMPOUNDS, METAL IONS OXIDATION-REDUCTION the rate law being ... 

COVALENT COMPOUNDS, METAL IONS OXIDATION-REDUCTION Steady-state treatment for L- leads to a rate law -d[02]/dr = fc2( i[Fe(acac)3])°- ... 

The cobaltous acetate reduction of tert-butyl hydroperoxide in acetic acid yields mainly ter/-butanol and oxygen the metal ion stays in the +2 oxidation state because of the reactivity of Co(III) towards hydroperoxides (p. 378) °. The rate law is... 

Figure 6. Reservoir sizes, residence times, and 5 Fe values for aqueous Fe(II), as calculated for DIR assuming first-order rate laws. Timescale arbitrarily set to 100 days. Calculations based on rate constant determined for a 23 day DIR experiment involving hydrous ferric oxide (HFO) by S. algae (Beard et al. 1999). The percent total reduction at 100 days is shown in the grey box on the lower right side of the lower diagrams, based on the value of k. Parts A-C assume a 2/ 1 ratio of 10, whereas parts D-F assume Bikjki ratio of 1000. As constrained by first-order rate laws, the proportion of the intermediate products Fe(III)-L, followed by Fe(II)-L, increase before substantial accumulation of the final Fe(II)aq product (Parts A and D). Tlie fraction of Fe(III)-L in the exchangeable pool of Fe (Fe(III)-L + Fe(II)-L + Fe(II)aq) decreases with time, primarily due to accumulation of the Fe(II)aq end product, where the rate of change is a function of the kjk ratio.

Monomer, polymer equilibria can be the basis of a genuine fractional order. Many reductions by dithionite S20 of oxidants (ox) to produce reductants (red) contain in the rate law a term which includes a square root dependence on the S204 concentration (often this is the... 

The aquated iron(III) ion is an oxidant. Reaction with reducing ligands probably proceeds through complexing. Rapid scan spectrophotometry of the Fe(III)-cysteine system shows a transient blue Fe(lII)-cysteine complex and formation of Fe(II) and cystine. The reduction of Fe(lII) by hydroquinone, in concentrated solution has been probed by stopped-flow linked to x-ray absorption spectrometry. The changing charge on the iron is thereby assessed. In the reaction of Fe(III) with a number of reducing transition metal ions M in acid, the rate law... 

Reduction of these PHMs took place in experiments containing both Fe(ll) and iron oxide minerals, under anoxic conditions. The transformation of PHMs by surface-bound Fe(ll) generally follows a pseudo-first-order kinetic rate law, expressed by... 

The rate law for the oxidation of [Ru(NH3)5(FlL)] + (HE = isonicotinamide) by I2 in acidic solution contains two terms, one depending on P2] and one depending on [I3 ] and [Ru complex]. An outer-sphere electron-transfer mechanism is proposed for each term. Reduction of [Ru (NFl3)5L] + (TIL = nicotinamide or isonicotinamide) to [Ru (NH3)5L]+ is accompanied by an isomerization from the amide-bonded L to pyridine-bonded FIL. Bromine oxidation of... 

Tris(diimine)ruthenium(III) complexes are significantly more oxidizing than the analogous complexes of both iron(III) and osmium(III). This correlates well with the observation that rates of reduction in base are also faster for the tris(diimine)ruthenium(III) complexes. The tris(l,10-phenanthroline)ruthenium(III) reduction is significantly faster than the tris(2,2 -bipyridine)ruthenium(III) reduction, and this may be the reason why it is only the latter reaction that has been investigated in detail (1, 2). This system is particularly complex, and the rate law given by Eq. (1) holds only for very small concentrations of ruthenium complex. In contrast to the irondll) systems, simple kinetics... 

Abiotic oxidation of sulfide by oxygen cannot supply sulfate at rates comparable to rates of sulfate reduction. Unless high concentrations of sulfide develop and the zone of oxidation is much greater than 1 cm, rates of chemical oxidation of sulfide by oxygen will be much less than 1 mmol/m2 per day (calculated from rates laws found in refs. 115-118). Such conditions can exist in stratified water columns in the Black Sea water column chemical oxidation rates may be as high as 10 mmol/m2 per day (84). However, in lakes in which sulfide is undetectable in the water column and oxygen disappears within millimeters of the sediment-water interface (e.g., 113), chemical oxidation of sulfide by oxygen is unlikely to be important. 

We assume the same rate law to generally apply when B is a one-electron reductant. The ability to prove the rate law for B=I stems from a rather low value of ket, Table I. Interest in the photochemical oxidation of I- to I - for energy storage purposes and the 10-fold difference in ket in EtOH vs. H2O prompted us to determine ket for B=I for several other solvents. Values of ket, %(I-0xdn) 311(1 E0(T,eCp2+/°)surf. in the vari°us solvents used are given in Table I. 

Estimation of rates for redox reactions in environmental systems requires that the problem be formulated in terms of specific oxidation and reduction half-reactions. In addition, we assume that the rate-limiting step of the transformation mechanism is bimolecular—that is, the slow step requires an encounter (collision) between the electron donor and electron acceptor. Under most conditions found in environmental systems, such reactions exhibit rate laws for the disappearance of a pollutant, P, that are first-order in concentration of P and first-order in the concentration of environmental oxidant or reductant, E,... 

In many cases, the concentration of the environmental oxidant or reductant is effectively constant over the time frame of interest, so Equation (21) can be simplified to a pseudo-first-order rate law... 

This rate law places two molecules of HBr in the transition state along with one molecule of Mn(III). This information was used to argue for HBr2" as a key intermediate and against the thermodynamically more energetic Br. Intramolecular oxidation of one of the two coordinated bromide ions followed by direct elimination of HBr2 appears to provide a convenient mechanism for one-electron oxidation of bromide by Mn(ffl). This is followed by rapid reduction of an additional equivalent of Mn(III) by HBr/ ... 

These laws (determined by Michael Faraday over a half century before the discovery of the electron) can now be shown to be simple consequences of the electrical nature of matter. In any electrolysis, an oxidation must occur at the anode to supply the electrons that leave this electrode. Also, a reduction must occur at the cathode removing electrons coming into the system from an outside source (battery or other DC source). By the principle of continuity of current, electrons must be discharged at the cathode at exactly the same rate at which they are supplied to the anode. By definition of the equivalent mass for oxidation-reduction reactions, the number of equivalents of electrode reaction must be proportional to the amount of charge transported into or out of the electrolytic cell. Further, the number of equivalents is equal to the number of moles of electrons transported in the circuit. The Faraday constant (F) is equal to the charge of one mole of electrons, as shown in this equation ... 

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