Association of Ferric Ions with Chloride, Bromide and Hydroxyl Ions (A Spectroscopic Study)

The kinetics of oxidation of uranium (IV) b y iron (111) in aqueous solutions of perchloric acid have been investigated a t four temperatures between 3.1°C. and 24.8"C. The reaction was followed by measurement of the amount of ferrous ion formed. For the conditions (Il+) = 0.1-1.0 iM, ionic strength = 1.02, (FelII) = 10-'-10-5 hf, and (UIV) = lO-'-lO-6 ill, the observed rate law is d(FeZ+)/dt = -2d(UIV) /dt -K1 and Kn are the first hydrolysis constants for Fe3+ and U4+, respectively, and K'and K" are pseudo rate constants. At 24.8"C., K' = 2.98 sec.-I, and K" = 10.6 mole liter-' sec-'. The corresponding temperature coefficients are AH' = 22.5 kcal./mole and AH" = 24.2 kcal./mole. The kinetics of the process are consistent with a mechanism which involves, as a rate-controlling step, electron transfer between hydrolyzed ions. INTRODUCTIONOxidation-reduction reactions between cations in aqueous solution have been thought until recently t o proceed very rapidly, and the term "instantaneous" has frequently been used to characterize their rates. However recent studies with radioactive tracers have shown that even such elementary reactions as electron transfer between iron (11) and iron (I 11), or between Ce (I 11) and Ce (IV), do not proceed instantaneously (18,8). Such results have reawakened interest in the whole subject of both isotopic and ilo~lisotopic electron transfer reactioils between cations in solution. Some kinetic information is now available relating to the reactioils between the nonisotopic pairs Sn (11) -Fe (111) (5), T1 (111) -Fe (11) (2), and N p (IV) -Fe (111) (10).Several review articles have recently appeared in which the current status of theory and experiment have been summarized for both isotopic and nonisotopic clcctron transfer reactioils (1,14,19). The proceedings of a symposium held in 1954 on this subject have also been published (12).As a further example of this type of reaction, the present paper describes experiments relating to the kinetics of the oxidatioil of uranium (IV) by iron (111) in aqueous acidic media. Perchloric acid and sodium perchlorate were chosen as the principal electrolytes, since amongst the common inorganic anions, perchlorate has the least tendency to form ion-pairs with cations. T o bring the rate of reaction into a range suitable for study, it has been necessary to work with solutions to 10-5 M in uranium (IV) and iron (111). T h e course of the reaction has been followecl by measurement of the rate of formatioil of iron (11). Such measurements, in conju~lction with a knowledge of the stoichiometry of the reaction, were sufficient to enable a kinetic analysis of the results to be made. I t would have been useful, as confirmatory evidence, lMa?zz~script

“…T h e principal hydrolyzed species present in solution are FeOH2+ and UOH3+ (15,17,13), which exist in equilibrium with the parent fully hydrated ions according to:…”

Section: Effect Of Hydrogen-ion Concentrationmentioning

confidence: 99%

“…The doubly hydrolyzed ion Fe(OH)2+ is also present a t very low concentration (8,9,16,17): Although there is no quantitative information available, it is reasonable to assume, by analogy with the chemistry of iron (111), that the ion U(OH)2++ also is present a t very low concentration:…”

Section: Effect Of Hydrogen-ion Concentrationmentioning

confidence: 99%

Kinetics of the Oxidation of Uranium (Iv) by Iron (Iii) in Aqueous Solutions of Perchloric Acid

Betts¹

1955

“…T o obtain values of K and E from optical densities somewhat similar methods are available, devised by Rabinowitch and Stocltmayer (19), and by Franlt and Oswalt (7). The relation of these two methods can be seen from the following, which provides a simpler derivation of Frank and Oswalt's equation.…”

Section: Discussion Of Resultsmentioning

confidence: 99%

Some Measurements on the Iron (111)–thiocyanate System in Aqueous Solution

Lister¹,

Rivington²

1955

A spectrophotometric study of the ferric thiocyanate system is reported. The temperature and ionicstrength of the solutions were carefully controlled, and suitable precautions were taken against fading. From the data, values of the equilibrium constant for the formation of FeSCN++ from simple ions, and of the extinction coefficients of FeSCN++ were obtained for different temperatures and ionic strengths, with results that differed somewhat from earlier values. The effect of varying acidity was also investigated. From more concentrated solutions, values for the equilibrium constants for the formation of Fe(SCN)Z+ and Fe(SCN)a were obtained; it was not necessary t o postulate higher complexes.Many investigations have been made of aqueous solutions containing ferric and thiocyanate ions, starting with the original work of Gladstone (9) just over a hundred years ago. However it was only comparatively recently that Moeller (15) and, independently, Bent and French (3) showed t h a t in dilute solution the color was due t o the ion FeSCN++. This put a new interpretation on the measurements made on this system, and as a result considerable further work was carried out (e.g. 6,7,10, 12,18). This work mostly made use of the red color a s a means of investigation, but some work was also done using solvent extractioil (14), or electrochemical methods (15).The present work was begun with the intention of using the red color of FeSCN++ as a method of determining the concentration of free ferric ions in solution, and hence of evaluating the equilibrium constants of the dissociation of other complexes formed by ferric ions with other anions. This type of investigation was made by Lanford and Iciehl (13) for the ferric phosphate system. Although earlier work covered a fair range of conditions, i t became apparent t h a t in order t o have an adequate picture of the system before adding other ions, i t would be desirable t o extend these measurements t o cover systematically a range of temperature, hydrogen ion concentration, and ionic strength. These factors were therefore carefully controlled; temperature control, for instance, was provided by using absorption cells completely immersed in a thermostatted tank. The fading which occurs in these solutions had also t o be talten into account, and it is believed conditions were found such t h a t fading was negligible. I t was also necessary t o talte into account the possible formation of the higher complexes, Fe(SCN)2+ and Fe(SCN)3; and it was in fact possible t o malte a moderately accurate estimate of the extent t o which these species were present, and so t o determine their dissociatioll constants. I n interpreting the experimental data, it is not necessary t o assume t h a t the extiilction coefficients a t various wave lengths were constant over a range of conditions; so that these experiments enable one t o find the effect of temperature and ionic strength on the extinction coefficients. The effect of hydrogen ion concentration was 'ilCa7tuscript received Jzote SO,...

“…For all phenolic ligands FeL has an absorption band in the visible spectral region so that, in principle, the position of equilibrium should be capable of measurement if the optical density and pH of the solutions containing iron(II1) and phenol are known. In practice, however, E, the molar extinction coefficient of FeL, is not directly obtainable, so that the experiments must be by no means a new problem, and a number of methods have been described for arranging data from experimental measurements in such a way as to extract values of E and K when dealing with the formation of weak complexes (2)(3)(4)(5)(6).…”

Section: Pi [Fe][hl] -K 'mentioning

confidence: 99%

Constraints on the determination of stability constants for metal complexes. II. The iron(III) phenolates

McBryde

1968

Stability constants have been measured spectrophoton~etrically for the 1 :1 co~nplexes of iron(1II) with phenol and five derivatives in background solutions having two different con~positions, 0.027 M NaC10, and 0.5 M KNO,. Conditional acidity constants for these phenols in the second of these media were also determined. The experin~ental results were treated in two different ways to obtain values of the equilibriunl constants and the molar extinction coefficients of the complexes. It is apparent that the equilibrium in these cases must be studied under conditions, such as low pH and low concentration of iron, which militate against good precision in the results.