KINETICS AND REACTION MECHANISM OF DECOMPOSITION OF HEXA AND OCTA COORDINATED CYANOGEN COMPLEXES BY METAL IONS

Abstract

The first systematic study on metal cyanides may be said to have began with the discovery of iron cyanides. Since then many basic and applied aspects related with iron cyanides and other metal cyanides have been comprehensively and critically examined. The various aspects covered are: the use of electrometric potentiometrie, polarographic and coulometric) methods in ascertaining their structure (1-5) ,study of their colloid chemical behaviour (6-10) (preparation and properties of their colloidal solutions, preparation of their gels, study of flow properties etc.) , use of radiochemical methods to establish their stoichiometry (11) ,X-ray analysis(12,13) preparation of their organic derivatives(14,15,16) , study of permeability of electrolytes through their membranes(17) . Interestingly enough until recently lesser attention was paid to study these compounds from the kinetic point of view inspite of the fact that the results of these studies can prove to be both of fundamental and applied importance. The first indication of the possibility investigating metal cyanides from the kinetic point of view came from the work of Pinter(18) who suggested that Hg++ ions in slightly acid solution catalyse the formation of Prussian blue from alkali hexacyanoferrate(II) . However,his suggestion remained almost obscure for about a decade and it was only in the early fifties that studies in this direction were taken up by Emschwiller in France and Asperger in England and Yugoslavia. They comprehensively studied the dissociation and decomposition of hexacyanoferrate (II) ion by H* and Hg++ ions and further extended these by irradiating the hexacyano ferrate solution by ultraviolet light. Emschwiller(19) initially started with the studies on the dissociation of hexacyanoferrate(II) ion in aqueous solution by H* ions. He observed that the rate of decomposition was first order in its varying concentration,but was zero order in the concentration of R* at high acidities. Further that the HCN formed during the dissociation retarded the decomposition at lower acidities. For example,0.01M K4Fe(CN) 6 dissociation reaction appeared homogenous in 1.0 to 5.ON and could be followed by determining Fe++ or by making use of the violet coloration given by nitrosobenzene with the reaction product. Enough evidence for stepwise dissociation first to Fe(CN) 5. Ha0 3~ and then to Fe++ could be obtained on the basis of these studies. The following mechanism for the reaction was put forward: Fe(CN) 64- + HaO-^eCCN) 5. HaO] &~ +CN" This slow reaction was followed by the capture of CN~ by H*. The activation energy of the reaction was 30.8 K. Cals. The effect of Hg ions on the aqueous solution of potassium hexacyanoferrate(II) was studied by U.V.light and Hg++ was studied by Asperger(20) . He observed that the equilibrium: Fe(CN) 64" + Ha0^=i Fe(CH) 5. Ha03~ + CN" in the dark although shifts completely towards the left the interaction between JFe(CN) 5. Had} 3~ and CN~ can be prevented by means of the relatively fast irreversible process: Ife(CK) 5. Ha0l?~ + C6H5N0 —^©(CH) 5. CeHsNOl 3~+ Ha0 The initial reaction velocity was found to depend on the concentration (5 xl0MM to 5 xlO"4^! ) of potassium hexacyano ferrate (I I) and also on the concentration of nitro so benzene but only up to a concentration of 1.4 xlO"aM. Addition of small amounts of HgCla (order 10~5 ) had the same influence as the ultraviolet light and the violet nitro so benzene complex was found to have the same spectra in both the cases. The velocity of the reaction of potassium hexacyano ferrate(II) and nitro so benzene in tne presence of Hg ions was pH dependent being maximum at pH 3.5. Furthermore on carrying out absorption experiments at this pH with different concentration of it J.J. Hg ions, they found that the catalytic action of Hg ions was so large that it could be detected at the concentration 10~7M . They also put forward the view that the reaction is strictly specific for Hg ions and a method can be developed for the determination of Hg ions in distilled water and that other metal cations did not interfere in the estimation. Based on this . a catalytic reaction they could also evolve method of the determination of mercury vapours in the atmosphere (21). The results of his studies can be summarised as follows: (i) The reaction takes place much faster in ultraviolet light then in the dark. (ii) The same violet product is formed in the presence of nitrosobenzene regardless of whether the reaction is catalysed by ultraviolet light or by Hg++ ions. (iii) Light below 400 mu caused the dissociation. (Iv) The catalytic action of mercuric ions must relate to the slowest process,i. e. , the decomposition of Fe(CN) 6*~ to Fe(CN)53~. (v) The product of reaction between Hg++ and Fe(CN) e4~is probably the intermediate,according to the scheme: Hg++ +(/e(CN)|4"^gQg5 £Fe(CN)5.HaOl3-+I^(^0" Hg(CN)+ + H+ ==* Hg++ + HCN and hence the catalytic action should be almost negligible in the basic medium. (vi) The effect of H4" ions is more complex than metal ions. In acid solution the following equilibria are possible: H* + ^e(CN) 3*^=5-HH"e(CrD §3~ H* + HfFe(CN) gf^Ha^e(CN) $*~ (vii) Smaller the number of protons in the complex ions, easier is the loss of cyanide ion. Consequently in highly acid solutions of K4Fe(CN) 6 the reaction velocity should decrease as is actually seen. The two effects of H4",suppression of the dissociation of H4Fe(CN) e and the regeneration of the catalyst, are in opposition and this must lead to optimum pH (3.5-4.0 ). (viii) The negative salt effect and the decrease in activation energy of the decomposition of Fe(CN) 64~ to Fe(CN) 53~ in presence of mercuric ions and that too for purely covalent mercury support the foregoing discussion. (ix) Ions having close affinity for the cyanide radical can also catalytically influence the reaction. The increasing order of catalytic action correspond to the arrangement of these metals in the periodic table Pt4+ < Au+++ < Hg++ < Hgt+ On the other hand Emschwiller(22) considers the decomposition of hexacyano ferrate (I I) as both catalytic and stoichiometric in nature and according to him both the reactions take place simultaneously. The reaction proceeds in the fo llowing two s tep s: LFe(CN) ^- + Hg++ ^Fe(CN) $>' +Hg(CN) + and Hg(CM)+ +[Fe(CN) g^if'e(CN) sf' + Hg(CN) a 6 In the presence of excess of hexacyano ferrate (II) the reaction is first order with respect to Hg ions while at lower concentration the reaction seem to be zero order with respect to Fe(CN) e4"". He further found that the Hg(CN) a formed hy stoichiometric reaction of mercuric salts also catalysed the decomposition of the hexacyanoferrate(II) . The rate constant decreases with time,probably due to the interaction of catalytically active Hg + and Hg(CN)+ with the decomposition product of hexacyanoferrate(II) . Asperger in his later communication(23) reported the photochemical decomposition of hexacyanoferrate(II) ion to aquopentacyanoferrate(II) ion and the analogous reaction involving release of a cyanide group from an octacyanomolybdate (IV) ion are both reversible in darkness. Talcing into consideration the reported observation of Mac Diarmid(24) that the pH of the solution which had been kept for some time in the light did not decrease in the dark(and who further observed that after short exposure to light the reversibility was much smaller then found by Asperger)-Asperger repeated his earlier observation/concluded that the reaction was reversible only when no appreciable decomposition of aquopentacyanoferrate( II) ion occured. However after exposing the 7 solution to ultra-violet light for several hours, the pH remained constant for a few hours but then decreased very slowly and returned after two days to almost its initial value. Under such conditions the decrease of pH did not mean that the reaction was reversible since a precipitate of Fe(OH) 3 was also formed. The same happened with the solutions of hexacyano ferrate (II) ion was exposed briefly in a quartz vessel to the U.V.light of a mercury quartz lamp or to sunlight for a few minutes only. Aqueous solution of octacyanomolybdate(IV) ion, exposed to U. V. light,decomposes readily and free cyanide ion are formed. The pH of irridiated solutions will, therefore,increase. As pointed out by Adamson and Sporer (25) a complicated reaction sequence is clearly involved, since the original yellow solutions develops a red colour which with prolonged irridation changes to green and finally to pale blue. The mechanism involves the release of cyanide ions in the first stage when the conditions were mild enough. This reaction is reversible in the dark. This was not the case with the irridiated solutions which acquired a pale blue colour. The pH of these blue solutions remained constant for some time but began to decrease slowly after a couple of days 8 probably because of the formation of basic cyanides.