Hexacyanoferrate ions reactions

The oxidative behaviour of glycolaldehyde towards hexacyanoferrate(III) in alkaline media has been investigated and a mechanism proposed, which involves an intermediate alkoxide ion. Reactions of tetranitromethane with the luminol and luminol-peroxide radical anions have been shown to contribute substantially to the tetranitromethane reduction in luminol oxidation with hexacyanoferrate(III) in aerated aqueous alkali solutions. The retarding effect of crown ethers on the oxidation of triethylamine by hexacyanoferrate(III) ion has been noted. The influence of ionic strength on the rate constant of oxidation of ascorbic acid by hexacyanofer-rate(III) in acidic media has been investigated. The oxidations of CH2=CHX (where X = CN, CONH2, and C02 ) by alkaline hexacyanoferrate(III) to diols have been studied. ... 

Copper(II) ions in aqueous solution are readily obtained from any copper-containing material. The reactions with (a) alkali (p. 430), (b) concentrated ammonia (p 413) and (c) hydrogen sulphide (p. 413) provide satisfactory tests for aqueous copper(II) ions. A further test is to add a hexacyanoferrate(II) (usually as the potassium salt) when a chocolate-brown precipitate of copper(II) hexacyanoferrate(II) is obtained ... 

Hexa.cya.no Complexes. Ferrocyanide [13408-63 ] (hexakiscyanoferrate-(4—)), (Fe(CN) ) , is formed by reaction of iron(II) salts with excess aqueous cyanide. The reaction results in the release of 360 kJ/mol (86 kcal/mol) of heat. The thermodynamic stabiUty of the anion accounts for the success of the original method of synthesis, fusing nitrogenous animal residues (blood, horn, hides, etc) with iron and potassium carbonate. Chemical or electrolytic oxidation of the complex ion affords ferricyanide [13408-62-3] (hexakiscyanoferrate(3—)), [Fe(CN)g] , which has a formation constant that is larger by a factor of 10. However, hexakiscyanoferrate(3—) caimot be prepared by direct reaction of iron(III) and cyanide because significant amounts of iron(III) hydroxide also form. Hexacyanoferrate(4—) is quite inert and is nontoxic. In contrast, hexacyanoferrate(3—) is toxic because it is more labile and cyanide dissociates readily. Both complexes Hberate HCN upon addition of acids. 

Because of the time and expense involved, biological assays are used primarily for research purposes. The first chemical method for assaying L-ascorbic acid was the titration with 2,6-dichlorophenolindophenol solution (76). This method is not appHcable in the presence of a variety of interfering substances, eg, reduced metal ions, sulfites, tannins, or colored dyes. This 2,6-dichlorophenolindophenol method and other chemical and physiochemical methods are based on the reducing character of L-ascorbic acid (77). Colorimetric reactions with metal ions as weU as other redox systems, eg, potassium hexacyanoferrate(III), methylene blue, chloramine, etc, have been used for the assay, but they are unspecific because of interferences from a large number of reducing substances contained in foods and natural products (78). These methods have been used extensively in fish research (79). A specific photometric method for the assay of vitamin C in biological samples is based on the oxidation of ascorbic acid to dehydroascorbic acid with 2,4-dinitrophenylhydrazine (80). In the microfluorometric method, ascorbic acid is oxidized to dehydroascorbic acid in the presence of charcoal. The oxidized form is reacted with o-phenylenediamine to produce a fluorescent compound that is detected with an excitation maximum of ca 350 nm and an emission maximum of ca 430 nm (81). 

The stability of complex ions varies within very wide limits. It is quantitatively expressed by means of the stability constant. The more stable the complex, the greater is the stability constant, i.e. the smaller is the tendency of the complex ion to dissociate into its constituent ions. When the complex ion is very stable, e.g. the hexacyanoferrate(II) ion [Fe(CN)6]4", the ordinary ionic reactions of the components are not shown. 

In strongly acid solution the reaction proceeds from left to right, but is reversed in almost neutral solution. Oxidation also proceeds quantitatively in a slightly acid medium in the presence of a zinc salt. The very sparingly soluble potassium zinc hexacyanoferrate(II) is formed, and the hexacyanoferrate(II) ions are removed from the sphere of action ... 

Such free radicals may be stabilized by binding to proteins. Redox reactions may also occur between ionic species, for example the oxidation of reduced cytochrome c by hexacyanoferrate (ferricyanide) ions. 

When oxidized by iron(III) ions 4-aminoantipyrine reacts with phenols to yield colored quinonoid derivatives (cf. 4-aminoantipyrine — potassium hexacyanoferrate(III) reagent in Volume 1 a). It is an oxidative coupling based on the Emerson reaction. 

These possess the dangerous reactions of the CN ion, or cyano group. The hexacyanoferrate (III) anion also has oxidising properties. Finally, it is thought to produce an extremely unstable acid in certain conditions. These salts also produce hydrogen cyanide, which is highly toxic, in an acid medium (see p.334). 

The participation of cations in redox reactions of metal hexacyanoferrates provides a unique opportunity for the development of chemical sensors for non-electroactive ions. The development of sensors for thallium (Tl+) [15], cesium (Cs+) [34], and potassium (K+) [35, 36] pioneered analytical applications of metal hexacyanoferrates (Table 13.1). Later, a number of cationic analytes were enlarged, including ammonium (NH4+) [37], rubidium (Rb+) [38], and even other mono- and divalent cations [39], In most cases the electrochemical techniques used were potentiometry and amperometry either under constant potential or in cyclic voltammetric regime. More recently, sensors for silver [29] and arsenite [40] on the basis of transition metal hexacyanoferrates were proposed. An apparent list of sensors for non-electroactive ions is presented in Table 13.1. 

In conclusion, the unique properties of Prussian blue and other transition metal hexa-cyanoferrates, which are advantageous over existing materials concerning their analytical applications, should be mentioned. First, metal hexacyanoferrates provide the possibility to develop amperometric sensors for non-electroactive cations. In contrast to common smart materials , the sensitivity and selectivity of metal hexacyanoferrates to such ions is provided by thermodynamic background non-electroactive cations are entrapped in the films for charge compensation upon redox reactions. 

Studies of medium effects on hexacyanoferrate(II) reductions have included those of dioxygen,iodate, peroxodisulfate, - [Co(NH3)5(DMSO)] +, and [Co(en)2Br2]+. Rate constants for reaction with dioxygen depended strongly on the electron-donor properties of the organic cosolvent. Rate constants for reduction of peroxodisulfate in several binary aqueous media were analyzed into their ion association and subsequent electron transfer components. Rate constants for reduction of [Co(en)2Br2] in methanol water and dioxan water mixtures were analyzed by a variety of correlatory equations (dielectric constant Grunwald-Winstein Swain Kamlet-Taft). 

By comparison with G e, relatively few independent measurements of G( OH) have been made. In contrast to, only the relative change in G( OH) with time has been reliably measured by pulse radiolysis [51]. In practice, absolute values of G( OH) have been obtained from scavenger studies or by material balance (reaction (7)). Fig. 7 shows data for aerated solutions of formate ion [52] and hexacyanoferrate(II) [53] taken from Fig. 1 of Ref. 54. The data for formic acid, which were included by LaVerne and Pimblott [54], have been omitted here because they were obtained at low pH where the primary yields are different (see Section 3.4). The solid line shows the best fit obtained using Eqs. (16) and (17) and the broken line is the best fit when the term u[5]/2 is omitted from Eq. (17). The respective sets of parameters are a = 1.64 and 1.69 nsec, g( OH) = 2.53 and 2.50 molecules (100 eV) and G°( OH) = 4.48 and 4.86 molecules (100 eV) These values differ significantly from those obtained by LaVerne and Pimblott [54], which were a = 0.258 nsec, g(" OH) = 2.66 molecules (100 eV) and G°( OH) = 5.50 molecules (100 eV) The reason for the difference is that LaVerne and Pimblott [54] chose G°( OH) = 5.50 molecules (100 eV) ... 

More generally, the method of competition kinetics is used to determine H-atom rate constants. The hexacyanoferrate(III) ion is a suitable solute because reaction (39) can be followed from the decrease in absorbance at 420 nm due to Fe(CN)g (8420 = 104 m mol ). When a second solute is present so that reaction (40) competes with reaction (39), G(-Fe (CN)g ) is given by Eq. (41) ... 

This process can be described in terms of a heterogeneous reaction in which ferri-cyanide (or hexacyanoferrate(III)) ions, [Fe(CN)g], are formed. At the beginning of the voltammetric peak, the current is controlled by the kinetic of the electron transfer across the electrode/electrolyte barrier so that the current increases somewhat exponentially with the applied potential. The value of the current is controlled 150-200 mV after the voltammetric peak by the diffusion rate of ferrocyanide ions from the solution bulk toward the electrode surface. 

The cyanide ion, CN, is isoelectronic with carbon monoxide and has an extensive chemistry of reaction with transition metals (e.g. the formation of the hexacyanoferrate(III) ion, [Fe(CN)63 ] by reaction with iron(III) in solution) but, unlike CO, it shows a preference for the positive oxidation states of the elements. This is mainly because of its negative charge. 

To extract the iron(III) ions, pass 20 ml of a hot 2 M hydrochloric acid solution through the cation exchanger. Gather the solution flowing out from the column in a 250-ml flask. Use the reaction with potassium hexacyanoferrate(II) to check for the completeness of iron extraction. 

One group of NADH oxidants, which does not fit the proposed reaction scheme in Fig. 2.4 are the metal complexes. Examples of this type include nickel hexacyanoferrate deposited on porous nickel electrodes [29], gold electrodes modified with cobalt hexacyanoferrate films [30] and adsorbed l,10-phenanthroline-5,6-dione complexes of ruthenium and osmium [31]. It is unclear how these systems work and no mechanism has been proposed to date. It may be worth noting that dihydronicotinamide groups have been shown to reduce aldehydes in a non-enzymatic reaction when the reaction is catalysed by zinc, a metal ion [15]. In a reaction between 1,10-phenanthroline-2-carboxaldehyde and N-propyl-l,4-dihydronicotinamide, no reaction was seen in the absence of zinc but when added to the system, the aldehyde was reduced and the nicotinamide was oxidised. This implies that either coordination to, or close proximity of, the metal ion activates... 

We have already discussed several cases of fast Fe(III) oxidations which occur by a non-bonded electron-transfer mechanism (Tables 13 and 14). One case of a relatively slow reaction, involving the substitution-inert hexacyanoferrate(III) ion, is shown in Table 14 (entry no. 17) and clearly demonstrates the electron-transfer oxidizing properties of this species with respect to easily oxidized aliphatic amines. Whether the same mechanism holds for compounds more resistant to oxidation, such as methylnaphthalenes (Andrulis et al., 1966) remains to be seen (the estimated rate constant at 25°C is ca. 10-7 M l s-1). Generally, hexacyanoferrate(III) seems to be a good non-bonded electron-transfer reagent (for a review, see Rotermund, 1975). 

The reaction between hexacyanoferrate(III) (commonly called ferricyan-ide and abbreviated below as Feic) and iodide ions... 

Similarly, from the Tables 1.17 and 1.18 we can see, for example, that permanganate ions (in acid medium) can oxidize chloride, bromide, iodide, iron(II), and hexacyanoferrate(II) ions, also that iron(III) ions may oxidize arsenite or iodide ions but never chromium(III) or chloride ions etc. It must be emphasized that the standard potentials are to be used only as a rough guide the direction of a reaction will depend on the actual values of oxidation-reduction potentials. These, if the concentrations of the species are known, can be calculated easily by means of the Nernst equation. 

The hexacyanoferrate(II) ion being a complex ion does not give the typical reactions of iron(II) (cf. Sections 1.31 to 1.33). The iron present in such solutions may be detected by decomposing the complex ion by boiling the solution with concentrated sulphuric acid in a fume cupboard with good ventilation, when carbon monoxide gas is formed (together with hydrogen cyanide, if potassium cyanide is present in excess) ... 

Iron(III) ions cannot be detected in a solution of hexacyanoferrate(III) with the usual reactions. The complex has to be decomposed first by evaporating with concentrated sulphuric acid or by igniting a solid sample, as described with hexacyanoferrate(II) (cf. Section III.21, reaction 5). 

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