Redox reaction ferric/ferrous ions

Fig. 9.7. An electrochemical cell consisting of a redox electrode reaction of hydrated ferric-ferrous ions and the standard hydrogen electrode reaction.

Photoreactions that involve transition metal ions, complexes or compounds can usually be classified as (photo)redox (simultaneous oxidation and reduction) processes. A representative non-photoassisted catalytic system is Fenton s reagent that produces HO radicals on reaction of ferrous ions (Fe2 +) and hydrogen peroxide (Scheme 6.287a). Its photochemical counterpart is the photo-Fenton process,1527 in which ferric (Fe3 + ) complexes in aqueous solutions (absorbing over 300 nm) are reduced to ferrous ions (Scheme 6.287b). 

Eq. (6.1) describes the simplest case of a redox reaction. An example could be the electron transfer between ferric/ferrous ions and an electrode... 

Such reactions are know as redox reactions and, in this case, the potential developed by the electrode system depends on the tendency of ferrous ions to donate an electron compared with the tendency of the ferric ions to accept one. 

Electrode reactions can be classified into two groups one in which an electron transfer takes place across the electrode interface, such as ferric-ferrous redox reaction (Fet, + e = Fe ) and the other in which an ion transfer takes place across the electrode interface, such as iron dissolution-deposition reaction (Fe M = FeVq). Since electrons are Fermi particles in contrast to ions that obey the Boltzmann statistics as described in Chap. 1, the reaction kinetics of the two groups differ in their electrode reactions. 

There is no redox couple in solution at the start of the ferrous-ceric titration because the solution contains only Fe ". The oxidation of ferrous to ferric occurs as soon as an aliquot of ceric ions enter the solution to effect the redox reaction shown in equation (4.1). The bulk of the initially present ferrous ions remain, with the ferric products of the redox reaction residing in the same solution, i.e. a Fe " ", Fe + redox couple is formed. This couple has the electrode potential Epf + pg2+. 

Both iron (II) and iron (III) form complexes with mercaptoacetic acid, SRSH2 (5, 11). The ferrous complexes, Fe(II) (RS)2-2 and Fe(II) (OH) (RS) , are highly air-sensitive and are rapidly oxidized to the intense red ferric complex, Fe(III)OH(RS)2 2 (5). Under air-free conditions the color of this latter complex is observed to fade at moderate to fast rates because of a redox reaction in which the iron is reduced to the ferrous state and the mercaptoacetate is oxidized to the disulfide. Michaelis and Schubert (9) proposed that the catalysis takes place through the alternate oxidation and reduction of iron ions in a sequence similar to that just described, but Lamfrom and Nielsen (4) were able to show that under mildly acid conditions the rate of oxygen uptake of solutions containing iron and... 

One of these, electron transfer, actually occurs in the ideal definitional sense. It applies to the few overworked redox reactions where there is no adsorbed intermediate. The ion in a cathodic transfer is located in the interfacial region and receives an electron (ferric becomes ferrous) without the nucleus of the ion moving. Later (perhaps as much as 10-9 s later), a rearrangement of the hydration sheath completes itself because that for the newly produced ferrous ion in equilibrium differs (in equilibrium) substantially from that for the ferric. Now (even in the electron transfer case) the ion moves, but the definition remains intact because it moves after electron transfer. The amounts of such small movements (changes in the ion-solvent distance for Fe2+ and Fe3+ ions in equilibrium) are now known from EXAFS measurements. 

The free-radicals are generated by discharge of proton, peroxides and easily reducible compounds at the cathode according to Eq. (1—4). The radial-anion of monomer is obtained by both direct and indirect electron transfer process [Eq. (5—6)]. The indirect process means that the active oxidizing species is formed from the electrolytes by electrode reaction, followed by interaction with the monomer. An unstable monomer like a,a -2-trichloro-p-xylene is formed and polymerizes instantaneously [Eq. (7)]. The regeneration of ferrous ion from the pool of used up ferric ion in a redox system is electrolytically successful [Eq. (8)] and an... 

In redox electrodes an inert metal conductor acts as a source or sink for electrons. The components of the half-reaction are the two oxidation states of a constituent of the electrolytic phase. Examples of this type of system include the ferric/ferrous electrode where the active components are cations, the ferricyanide/ferrocyanide electrode where they are anionic complexes, the hydrogen electrode, the chlorine electrode, etc. In the gaseous electrodes equilibrium exists between electrons in the metal, ions in solution and dissolved gas molecules. For the half-reaction... 

It is important to note that the electrode potential is related to activity and not to concentration. This is because the partitioning equilibria are governed by the chemical (or electrochemical) potentials, which must be expressed in activities. The multiplier in front of the logarithmic term is known as the Nernst slope . At 25°C it has a value of 59.16mV/z/. Why did we switch from n to z when deriving the Nernst equation in thermodynamic terms Symbol n is typically used for the number of electrons, that is, for redox reactions, whereas symbol z describes the number of charges per ion. Symbol z is more appropriate when we talk about transfer of any charged species, especially ions across the interface, such as in ion-selective potentiometric sensors. For example, consider the redox reaction Fe3+ + e = Fe2+ at the Pt electrode. Here, the n = 1. However, if the ferric ion is transferred to the ion-selective membrane, z = 3 for the ferrous ion, z = 2. 

Because any potentiometric electrode system ultimately must have a redox couple (or an ion-exchange process in the case of membrane electrodes) for a meaningful response, the most common form of potentiometric electrode systems involves oxidation-reduction processes. Hence, to monitor the activity of ferric ion [iron(III)], an excess of ferrous iron [iron(II)] is added such that the concentration of this species remains constant to give a direct Nemstian response for the activity of iron(III). For such redox couples the most common electrode system has been the platinum electrode. This tradition has come about primarily because of the historic belief that the platinum electrode is totally inert and involves only the pure metal as a surface. However, during the past decade it has become evident that platinum electrodes are not as inert as long believed and that their potentiometric response is frequently dependent on the history of the surface and the extent of its activation. The evidence is convincing that platinum electrodes, and in all probability all metal electrodes, are covered with an oxide film that changes its characteristics with time. Nonetheless, the platinum electrode continues to enjoy wide popularity as an inert indicator of redox reactions and of the activities of the ions involved in such reactions. 

Let us consider an electronic transfer reaction of the redox couple of ferrous-ferric ions ... 

Sibmooh, N. Udomsangpetch, R. Kijjoa, A. Chantharaksri, U. Mankhetkorn, S. Redox reaction of artemisinin with ferrous and ferric ions in aqueous buffer. Chem. Pharm. Bull Tokyo, 2001, 49(12) 1541-1546. 

As pointed out previously the redox potential of the reduction of ferric to ferrous ion is +770 mv relative to the standard hydrogen electrode (SHE). This means that this reaction will be driven in the forward direction whenever ferric ion is present in a system whose overall redox potential is lower than +770 mv. 

In the presence of a ferric-ferrous redox reaction in the solution, the anodic hole-emitting GaAs dissolution of Reaction 22.53 can be coupled with the cathodic reduction of hydrated ferric ions, which injects holes in the valence band as follows ... 

R. L. S. Willis, Ferrous-Ferric Redox Reaction in the Presence of Sulfate Ion. Trans. Faraday... 

Reactions occurring in reductionroxidation systems furnish radicals at a Considerably lower energy expenditure than that of a decomposition of a peroxide or a hydroperoxide. In the following equations, a redox reaction is exemplified by the oxidatioii of a divalent ferrous ion to a trivalent ferriC ion ... 

Persulfate ion initiator (e.g., fromKjSjOg) reacts with a reducing agent such as a bisulfite ion (e.g., from NaHSOj) to produce radicals for redox initiation (Equations 2.5 and 2.6). Ferric ion may also be used as a source of radicals (Equation 2.7). Other redox reactions involve the use of alkyl hydroxides and a reducing agent such as ferrous ion (Equation 2.8). 

The hybrid TLM described above can be further extended to the adsorption of metal ions on metal oxide, hydroxide, or sulfide surfaces. As an example, the adsorption of ferrous ions on sphalerite, which causes unintentional activation of sphalerite [39,40], is examined. Additional species to be considered include FePH and Fe(OH)2, while Fe(OH)3 and Fe(OH)4 are not considered because of their relatively low concentrations in the pH range of current interest (pH > 8). It has been noted that the ferrous species on the sphalerite surface oxidizes to the corresponding ferric species. Therefore, the redox reaction on a sphalerite surface needs to be considered by including the corresponding ferric hydroxy species in surface complexing reactions. The reactions considered in addition to those shown in Table 1 are given in Table 5. 

PD is an aqueous solution containing silver ions, a ferrous/ferric redox (reduction/oxidation) system, a buffer (citric acid), and a cationic surfactant (generally -dodecylamine acetate). The ferrous (Fe ) ions in solution reduce the silver (Ag +) ions to silver metal (Ag°), with ferric (Fe + ) ions being present to hold back the reaction (eqn [1]) ... 

Whether lipid oxidation in muscle foods is catalysed by the iron redox cycle or by formation of the ferryl ions is not clear. However, ferrous ions react with lipid hydroperoxides much faster than with hydrogen peroxide. As shown above, if the reaction of metmyoglobin with hydroperoxides produces ferryl radicals capable of initiating lipid oxidation, it is necessary to prevent the formation of metmyoglobin or methemoglobin. At acidic pH, ferric myoglobin can initiate lipid oxidation in the presence of lipid hydroperoxides. 

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