Ferrocyanide catalysis

Srikantan and Rao (79) reported that the ferrocyanide catalysis is not first order in peroxide as found by Kistiakowsky but much more complicated. No details of the experiments were given. They also studied the ferricyanide-peroxide reaction and found that in the dark there was a pronounced induction period after which the decomposition rate was first order in peroxide concentration. Illuminating for five minutes in bright sunlight removed this induction period but the subsequent first order rate is somewhat less than the original dark rate. They suggest that the induction period was the time taken to build up on... 

Prussian blue (PB ferric ferrocyanide, or iron(III) hexacyanoferrate(II)) was first made by Diesbach in Berlin in 1704.88 It is extensively used as a pigment in the formulation of paints, lacquers, and printing inks.89,90 Since the first report91 in 1978 of the electrochemistry of PB films, numerous studies concerning the electrochemistry of PB and related analogs have been made,92 with proposed applications in electrochromism1 and electrochemical sensing and catalysis 93... 

The electrochemical rate constants for hydrogen peroxide reduction have been found to be dependent on the amount of Prussian blue deposited, confirming that H202 penetrates the films, and the inner layers of the polycrystal take part in the catalysis. For 4-6 nmol cm 2 of Prussian blue the electrochemical rate constant exceeds 0.01cm s-1 [12], which corresponds to the bi-molecular rate constant of kcat = 3 X 103 L mol 1s 1 [114], The rate constant of hydrogen peroxide reduction by ferrocyanide catalyzed by enzyme peroxidase was 2 X 104 L mol 1 s 1 [116]. Thus, the activity of the natural enzyme peroxidase is of a similar order of magnitude as the catalytic activity of our Prussian blue-based electrocatalyst. Due to the high catalytic activity and selectivity, which are comparable with biocatalysis, we were able to denote the specially deposited Prussian blue as an artificial peroxidase [114, 117]. 

The above compensating reactions are attractive because of the success of similar schemes in the halide catalysis, but proof in this case is more difficult. Thus it was possible to show in the halide systems that halogen and halide are present simultaneously. Evidence for the presence of ferrous ion in the ferric catalysis would support a similar interpretation. Manchot and Lehmann (44) claimed to have proved that ferrous ion is formed from ferric ion in the presence of peroxide since the addition of <, < -dipyridyl to the mixture resulted in the slow formation of the red ferrous tris-dipyridyl ion Fe(Dipy)3++. However, later work (65,66), which will be discussed when these systems are considered in more detail (IV,6), indicates that the ferrous complex ion may be formed by reduction not of the ferric ion, but of a ferric dipyridyl complex. Similar conclusions on the presence of ferrous ion were drawn by Simon and Haufe (67) from the observation that on addition of ferri-cyanide to the system Prussian blue is formed. This again is ambiguous, since peroxide is known to reduce ferricyanide to ferrocyanide and the latter with ferric ion will of course give Prussian blue (53). 

The ready reversibility of the ferrocyanide-ferricyanide redox system makes it a potential catalyst for the decomposition of hydrogen peroxide by the mechanism of compensating oxidation-reduction reactions. Moreover, the well-known facts that in acid solution ferrocyanide is oxidized to ferricyanide, whereas in alkaline solution the reverse reduction occurs, seem a good indication that at suitable pH s both reactions might occur to give catalytic decomposition. But from the investigations to date it would appear doubtful whether any such catalysis occurs to a measurable extent, and that what seems to be ready reactions of ferro- and ferricyanides are in fact those of partial hydrolysis products of these ions in which water molecules replace the cyanide ions in the coordination shell. 

Lai and Singhal (86) have examined the catalysis by ferricyanide. Using much smaller concentrations (M/300) than those of Srikantan and Rao they find that the dark rate is about one hundred times slower than with the same concentration of ferrocyanide, and that mixtures of the two salts give about the same rates as ferrocyanide alone. Initial irradiation of ferricyanide alone gives an increase of about tenfold over the ferricyanide dark rate, but if new ferrocyanide is added to this solution the rate rises to about ten times that of the ferrocyanide dark rate. These observations are explained in terms of the reactions given above. 

Thus ferricyanide hydrolyzes photochemically to aquopentacyanoferrate which is a slightly better catalyst for decomposition than ferricyanide. On addition of ferrocyanide the ferrate is reduced to the ferrite which is more active. Some confirmation of these reactions is given by the observations that addition of the ferrate to ferricyanide increases the dark rate, and also that cyanide, nitrate, and nitrosobenzene inhibit this as well as the catalysis induced by irradiation of ferricyanide. 

The peroxidatic activity of HRP is the catalysis of reactions by hydrogen peroxide and certain other oxidizing agents (i) summarized in the following reaction scheme for H2O2 and the substrate ferrocyanide. 

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