Ferrous-ferric system

The standard potential of the indicator system is not known exactly, but experiments have shown that in not too strongly acid solutions the sharp color change from colorless to violet, with green as a possible intermediate, occurs at a potential of about — 0.75 volt. The standard potential of the ferrous-ferric system is 0.78 whereas that of the di-chromate-chromic ion system in an acid medium is approximately — 1.2 volt hence a suitable oxidation-reduction indicator might be expected to have a standard potential of about — 0.95 volt. It would thus appear that diphenylamine would not be satisfactory for the titration of ferrous ions by acid dichromate, and this is actually true if a simple ferrous salt is employed. In actual practice, for titration purposes, phosphoric acid or a fluoride is added to the solution these substances form complex ions with the ferric ions with the result that the effective standard potential of the ferrous-ferric system is lowered (numerically) to about — 0.5 volt. The change of potential at the end-point of the titration is thus from about — 0.6 to — 1.1 volt, and hence diphenylamine, changing color in the vicinity of — 0.75 volt, is a satisfactory indicator. 

In the presence of chloride ion, these reactions lead to the formation of chlorine atoms as well as azide radicals and the former react preferentially with organic (carbon) free radicals. This is demonstrated by the reaction of cyclohexanone peroxide with the ferrous-ferric system in presence of both azide and chloride ions to give exclusively -chlorocaproic acid (equation 111). This may be exploited syntheti-... 

Luther and Inglis (10) used the ferrous-ferric system to measure ozone at pH 2.0. An excess of ferrous ion is added to the ozone solution and the excess ferrous ion is titrated with permanganate. The author found this technique, also used by David (5) and Kawamura (7), to be highly precise and to give values for the ozone which agree with the equation... 

If the driving potential for the ozone against bromide at pH 2.0 is inadequate or more time is needed for the reaction, a slight excess of ozone may be present when the iodide is added for the titration. When chloride is added to the ozone solution at pH 2.0, the ozone odor persists in a closed reaction vessel for an hour, as discussed by Yeatts and Taube 16). After the chloride has been added at pH 0.0, there is a greater excess of iodine than that calculated from the ferrous-ferric system ozone may have been present at the time of the iodine addition. 

Apart from the qualitative evidence mentioned above which is in favor of a ferric ferrous change, it has been shown recently (8) that an expression which is, under certain conditions identical with Andersen s empirical Eq. (VIII) can actually be obtained from the free radical mechanism discussed above by taking into account the six Eqs. (1), (2), (4), (5), (6), and (7). Treating these equations for the stationary state of OH, HO2, and of the ferrous/ferric system one obtains ... 

An estimate of the free energy required to equalise the two bond lengths at 2.09 A in the ferrous-ferric system is around 30 kJ mol. Comparison with the observed activation energy of about 70 kJ mol suggests that, while the bonding changes are responsible for a major contribution, there are other factors as well. The main one is the reorientation of solvent molecules, to which we now turn. 

The last reaction cited above as shown is very effectively catalyzed by bacterial action but is very slow chemically by recycling the spent ferrous liquors and regenerating ferric iron bacterially, the amount of iron which must be derived from pyrite oxidation is limited to that needed to make up losses from the system, principally in the uranium product stream. This is important if the slow step in the overall process is the oxidation of pyrite. The situation is different in the case of bacterial leaching of copper sulfides where all the sulfide must be attacked to obtain copper with a high efficiency. A fourth reaction which may occur is the hydrolysis of ferric sulfate in solution, thus regenerating more sulfuric acid the ferrous-ferric oxidation consumes acid. 

After more than two decades of investigation, it is pertinent to ask to what extent the predictions of this type of theory, and its agreement with experiment, have changed. The system which has received most study is the aquo ferrous-ferric outer-sphere symmetrical exchange ... 

Merz and Waters (1949) showed that oxidation of organic compounds by Fenton s reagent could proceed by chain as well as non-chain mechanisms, which was later confirmed by Ingles (1972). Kremer (1962) studied the effect of ferric ions on hydrogen peroxide decomposition for Fenton s reagent. It was confirmed that once ferric ions are produced the ferric-ferric system is catalytic in nature, which accounts for relatively constant concentration of ferrous ion in solutions. 

T. P. Hoar (38a) has suggested that the different behavior of ferric and ceric ions in silver dissolution may be considered from a purely electrochemical viewpoint. The anodic and cathodic potentials are very close in the ferric system their polarization curves meet at rather small values of the corrosion current, which does not require that all ferric ion near the metal surface be reduced to ferrous. On the other hand the potential of the ceric-cerous couple is about 0.8 V more noble, and complete reduction at the interface (or complete concentration polarization with respect to ceric ion) is necessary to lower this potential to the value of the Ag-Ag+ couple. 

This low value of the equilibrium constant means that when equilibrium is attained in the ferrous-ferric and stannous-stannic mixture, the concentrations (activities) of ferric and stannous ions must be negligibly small in comparison with those of the ferrous and stannic ions. In other words, when these two systems are mixed, reaction occurs so that the ferric ions are virtually completely reduced to ferrous ions while the stannous are oxidized to stannic ions. This fact is utilized in analytical work for the reduction of ferric to ferrous ions prior to the estimation of the latter by means of dichromate. 

Minisci, Galli and co-workers have studied a variety of radical reactions which result in the formation of organic azides. These processes commonly involve the interaction of an organic peroxide, an alkene, and azide ion in the presence of a ferrous-ferric redox system. The initial step is the reduction of the peroxide by Fe " " to form a free alkoxy radical (Fenton reaction, equation 104). 

However, although the Haber-Willstatter chain reactions have been assumed to occur in certain catalytic systems, notably the ferrous-ferric ion system (4), more recent evidence to be described subsequently, does not support this assumption. On the other hand, such reactions appear to offer the most plausible explanation for the photochemical decomposition of hydrogen peroxide, although even here a satisfactory analysis of the kinetics has yet to be made. 

The general Haber-Weiss mechanism involves an intermediate reduction oxidation of the ferric-ferrous ion system. This has been accepted by a number of authors and Simon et al. (15) have shown that catalytic decomposition of hydrogen peroxide by ferric salts only takes place under conditions where the formation of ferrous ions can be demonstrated. There is, therefore, no experimental basis for a mechanism without the intermediate reduction of the ferric ions. 

The excess free energy associated with small platinum particles can be measured by the potential at which they are oxidised to Pt in the presence of chloride ion. An auxiliary redox system (Fe +/Fe +) had to be used, and the platinum particles then acted as a microelectrode, taking the reversible potential of the ferrous-ferric equilibrium, which was calculated by the Nemst equation. Use of different concentrations of ferric ion allowed the potential at which platinum atoms were oxidised to be determined, and from the equation... 

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]) ... 

Baron " and by Clark " in their studies on the potentials of the hemo-chromogens. The hemochromogens are thermodynamically reversible univalent redox systems, that is, they require a negligible activation energy to accept an electron in the oxidized state and to release an electron in the reduced state. The redox level of the iron in the heme may be varied by varying the N compounds that complex with the heme, and one may obtain E, values at pH 7 (Table III) as negative as —0.183 in the case of the ferrous-ferric protoporphyrin.2CN system and as high as... 

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