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

Abstract

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,...

The Cysteine Desulfhydrase CdsH Is Conditionally Required for Sulfur Mobilization to the Thiamine Thiazole in Salmonella enterica

“…The absorbance values for the buffer-only control (including KCN and the Sorbo reagent) were subtracted from the experimental values. The concentration of ferric thiocyanate was calculated using the following formula: ε 460 ϭ 4,700 M Ϫ1 cm Ϫ1 (43).…”

The Effect of Ionic Strength and Pressure on the Complex Formation and Hydrolysis Equilibria in the Aquated Iron(III)‐Thiocyanate System

“…It is encouraging to note that our value for & is of the same order as the corresponding value reported for Fe(H20)&I2+ [ 151. Many investigators have studied the formation of Fe(H20)5SCN2+ without considering the effect of the hydrolysis constant K4 [13,14,[16][17][18].…”

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

“…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).…”

“…[4] it was possible to estimate E. This could be applied in the left-hand side of eq. [5] to derive an improved value of EK, and the whole cycle repeated until constant values of E were obtained. Different combinations of C, and C,, could be taken to obtain independent evaluations of EK.…”