Hexacyanoferrate complex

Such cyanide complexes are also known for several other metals. All the fer-rocyanide complexes may be considered as the salts of ferrocyanic acid H4Fe(CN)e and ferricyanide complexes are that of ferricyanic acid, H3Fe(CN)e. The iron-cyanide complexes of alkali and alkaline-earth metals are water soluble. These metals form yellow and ruby-red salts with ferro-cyanide and ferricyanide complex anions, respectively. A few of the hexa-cyanoferrate salts have found major commercial applications. Probably, the most important among them is ferric ferrocyanide, FeFe(CN)e, also known as Prussian blue. The names, formulas and the CAS registry numbers of some hexacyanoferrate complexes are given below. Prussian blue and a few other important complexes of this broad class of substances are noted briefly in the following sections ... 

An example of a rapid self-exchange reaction in which the change in M-L distances is particularly small is that involving hexacyanoferrate complexes as shown in Equation (5.42) ... 

Manganese(II) can be titrated directly to Mn(III) using hexacyanoferrate(III) as the oxidant. Alternatively, Mn(III), prepared by oxidation of the Mn(II)-EDTA complex with lead dioxide, can be determined by titration with standard iron(II) sulfate. 

Probably the most extensively applied masking agent is cyanide ion. In alkaline solution, cyanide forms strong cyano complexes with the following ions and masks their action toward EDTA Ag, Cd, Co(ll), Cu(ll), Fe(ll), Hg(ll), Ni, Pd(ll), Pt(ll), Tl(lll), and Zn. The alkaline earths, Mn(ll), Pb, and the rare earths are virtually unaffected hence, these latter ions may be titrated with EDTA with the former ions masked by cyanide. Iron(lll) is also masked by cyanide. However, as the hexacy-anoferrate(lll) ion oxidizes many indicators, ascorbic acid is added to form hexacyanoferrate(ll) ion. Moreover, since the addition of cyanide to an acidic solution results in the formation of deadly... 

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. 

Exciting developments have occurred in the coordination chemistry of the alkali metals during the last few years that have completely rejuvenated what appeared to be a largely predictable and worked-out area of chemistry. Conventional beliefs had reinforced the predominant impression of very weak coordinating ability, and had rationalized this in terms of the relatively large size and low charge of the cations M+. On this view, stability of coordination complexes should diminish in the sequence Li>Na>K>Rb> Cs, and this is frequently observed, though the reverse sequence is also known for the formation constants of, for example, the weak complexes with sulfate, peroxosulfate, thiosulfate and the hexacyanoferrates in aqueous solutions. 

The actual catalyst is a complex formed from osmium tetroxide and a chiral ligand, e.g. dihydroquinine (DHQ) 9, dihydroquinidine (DHQD), Zj -dihydroqui-nine-phthalazine 10 or the respective dihydroquinidine derivative. The expensive and toxic osmium tetroxide is employed in small amounts only, together with a less expensive co-oxidant, e.g. potassium hexacyanoferrate(lll), which is used in stoichiometric quantities. The chiral ligand is also required in small amounts only. For the bench chemist, the procedure for the asymmetric fihydroxylation has been simplified with commercially available mixtures of reagents, e.g. AD-mix-a or AD-mix-/3, ° containing the appropriate cinchona alkaloid derivative ... 

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. 

Trace metals have to be removed, notably manganese, ferrous ions and zinc. This is often accomplished using the compound potassium hexacyanoferrate which predpitates or complexes the metals and, in excess, acts to inhibit growth and indirectly promotes dtric add production. The amount of potassium hexacyanoferrate required is variable depending on the nature of the ion content of the carbon source. 

The Lewis bases attached to the central metal atom or ion in a d-metal complex are known as ligands they can be either ions or molecules. An example of an ionic ligand is the cyanide ion. In the hexacyanoferrate(II) ion, [Fe(CN)6]4, the CN- ions provide the electron pairs that form bonds to the Lewis acid Fe2+. In the neutral complex Ni(CO)4, the Ni atom acts as the Lewis acid and the ligands are the CO molecules. 

FIGURE 16.16 When potassium cyanide is added to a solution of iron(ll) sulfate, the cyanide ions replace the H.O ligands of the [Fe(H20), - + complex (left and produce a new complex, the hexacyanoferrate(ll) ion, Fe(CN)(l 4 (right). The blue color is due to the polymeric compound called Prussian blue, which forms from the cyanoferrate ion. 

The richness of coordination chemistry is enhanced by the variety of shapes that complexes can adopt. The most common complexes have coordination number 6. Almost all these species have their ligands at the vertices of a regular octahedron, with the metal ion at the center, and are called octahedral complexes (1). An example of an octahedral complex is the hexacyanoferrate(ll) ion, [Fe(CN)f, 4. ... 

When it is heated above 196°C, the mixture of this compound with chromium trioxide combusts. When it is submitted to friction or impact, the same mixture detonates violently. The same happens if it is heated with sodium nitrite. Nitrite gives rise to a detonation with potassium hexacyanoferrate (II) too. The dangerous site of these complex anions is the cyano group. 

In the complex [Co(NH3)6]Cl3, the cation is [Co(NH3)6]3+, and it is named first. The coordinated ammonia molecules are named as ammine, with the number of them being indicated by the prefix hexa. Therefore, the name for the compound is hexaamminecobalt(III) chloride. There are no spaces in the name of the cation. [Co(NH3)5C1]C12 has five NH3 molecules and one CN coordinated to Co3+. Following the rules just listed leads to the name pentaamminechlorocobalt(III) chloride. Potassium hexacyanoferrate(III) is K3[Fe(CN)6j. Reinecke s salt, NH4[Cr(NCS)4(NH3)2], would be named as ammonium diamminetetrathiocyanatochro mate (III). In Magnus s green salt, [Pt(NH3)4][PtCl4], both cation and anion are complexes. The name of the complex is tetraammineplatinum(II) tetrachloroplatinate(II). The compound [Co(en)3](N03)3 is named as tris(ethylenediamine)cobalt(III) nitrate. 

Unfortunately, many compounds contain bonds that are a mixture of ionic and covalent. In such a case, a formal charge as written is unlikely to represent the actual number of charges gained or lost. For example, the complex ferrocyanide anion [Fe(CN)6]4- is prepared from aqueous Fe2+, but the central iron atom in the complex definitely does not bear a +2 charge (in fact, the charge is likely to be nearer +1.5). Therefore, we employ the concept of oxidation number. Oxidation numbers are cited with Roman numbers, so the oxidation number of the iron atom in the ferrocyanide complex is +11. The IUPAC name for the complex requires the oxidation number we call it hexacyanoferrate (II). 

Aniline black (Cl Oxidation Base 1) is a complex polymeric phenazine that can be produced on cotton fabric by impregnation with aniline hydrochloride and suitable inorganic oxidants, such as sodium chlorate, ammonium vanadate and copper hexacyanoferrate(II). Aniline black is also made directly for use as a pigment (Cl Pigment Black 1). 

Iron(III) very readily forms complexes, which are commonly 6-coordinate and octahedral. The pale violet hexaaquo-ion [Fe(H20)6]3+ is only found as such in a few solid hydrated salts (or in their acidified solutions), for example Fe2(S04)3.9H20. Fe(C104)3.10H20. In many other salts, the anion may form a complex with the iron(III) and produce a consequent colour change, for example iron(III) chloride hydrate or solution, p. 394. Stable anionic complexes are formed with a number of ions, for example with ethanedioate (oxalate), C204, and cyanide. The redox potential of the ironll ironlll system is altered by complex formation with each of these ligands indeed, the hexacyanoferrate(III) ion, [Fe(CN)6]3. is most readily obtained by oxidation of the corresponding iron(II) complex, because... 

Complex iron(III) salts are frequently used in oxidative arene coupling reactions and quinone formation and tetra-n-butylammonium hexacyanoferrate(III) has several advantages in it use over more conventional oxidative procedures. When used as the dihydrogen salt, Bu4N[H2Fe(CN)6], it oxidizes 2,6-di-z-buty 1-4-methylphenol (1) to the coupled diarylethane (2), or aryl ethers (3) and (4) (Scheme 10.4), depending on the solvent. It is noteworthy that no oxidation occurs even after two days with the tris-ammonium salt. 

There have been two books devoted to the chemistry of iron, " and many reviews devoted to various aspects of its coordination chemistry, including structures and photochemistry (iron(III)). Iron complexes appear in a multi-author volume on the history of coordination chemistry, but there is disappointingly little about iron—just a brief mention of hexacyanoferrates in connection with pigments—in an otherwise excellent overview of the history of chemistry. ... 

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