Iron forms complex ions

One of the best oxidation-reduction indicators is the 1,10-phenanthroline-iron(II) complex. The base 1,10-phenanthroline combines readily in solution with iron(II) salts in the molecular ratio 3 base l iron(II) ion forming the intensely red l,10-phenanthroline-iron(II) complex ion with strong oxidising agents the iron(III) complex ion is formed, which has a pale blue colour. The colour change is a very striking one ... 

The chromous salts, derived from the oxide CrO, arc analogous to the salts of divalent vanadium, manganese, and iron. This is seen in the isomorphism of the sulphates of the type R" SOj-THgO. The stability of such salts increases in the order of the atomic number of the metal. The chief basic oxide of chromium is the sesquioxidc CraO, which is closely allied to ferric oxide, and, like the latter, resembles aluminium oxide. The hydroxide, Cr(OH)3, with bases yields chromites analogous to, but less stable than, the aluminates. Chromic sulphate enters into the formation of alums. The chromic salts are very stable, but in the trivaJent condition the metal shows a marked tendency to form complex ions, both anions and cations thus it resembles iron in producing complex cyanides, whilst it also yields compounds similar to the cobaltamines. 

A similar situation appUes to Cd assumed to be coated with a film of Cd(OH)2 (solubility produet = 2 X The caleulated value of (fn is -0.54, which is noble to iron, contrary to an observed galvanie potential less noble than iron (see Fig. 3.3, Section 3.8). The observed more-active galvanie potential of cadmium is plausibly explained by the known tendency of <2d to form complex ions that lower Cd activity below the value corresponding to saturated <2d(OH)2. 

These considerations are important to the action of iron because of its ability to form complex ions with hydroxyl. At low pH, hydrolysis is incomplete and Fe(OH) " and Fe(OH)J exist. At high pH, ferric ion forms Fe(OH)J (or FeOj-2H20), and the solubility of iron increases. The solubility is at a minimum between pH 5 and pH 10. This range includes the normal pH at the membrane interface and explains the lack of penetration by iron. The accumulation of Fe(OH)3 can raise the cell voltage but does not aifect the current efficiency. When the pH profile shifts to aflow the interfacial value to drop below five, ions can penetrate the membrane and deposit within the membrane but closer to the anode side. 

Finally, we should briefly discuss the rate of steady state dissolution of a compact corrosion product in aqueous solution. In order to calculate this stationary corrosion rate, it is necessary to know the rate-determining step for the removal of metal ions from the corrosion product into the electrolytic solution. It has been found that the stationary corrosion rate of a passivated layer on iron between 0.7 and 3.9 in an aqueous H2SO4 solution does not depend upon the stirring rate, nor upon the Fe /Fe ratio in the solution, nor, most importantly, does it depend upon the electrode potential h [34]. From these observations it has been concluded that the rate-controlling step is a chemical reaction between the iron ions, which are leaving the corrosion product, and the solvent ions to form complex ions immediately next to the surface. Thus, in the case of iron under the given conditions, this step determines the overall corrosion rate and thus it also determines the stationary thickness of the passivating corrosion product. 

Chlorodicarbonyl(tetraphenylcyclobutadiene)cobalt loses carbon monoxide as well as chlorine when stirred with trifluoromethylthiosilver, to give the sulphur-bridged dinuclear complex [Ph4C4Co(CO)(S CF3)]2, contrasting with the formation of the mononuclear iron complex C5H5Fe(CO)2(S-CF3). Cobalt(ii) salts form complex ions [Co(HHIB)2] with hexafluoro-hydroxyisobutyric acid, which are not oxidized to cobalt(iii) under the reaction conditions. ... 

Iron can assume the oxidation states+2, +3, and +6, the last being rare, and represented by only a few compounds, such as potassium ferrate, KaFeOj. The oxidation states +2 and +3 correspond to the ferrous ion, Fe ", and ferric ion, Fe, respectively. The ferrous ion has six electrons in the incomplete 2>d subshell, and the ferric ion has five electrons in this subshell. The magnetic properties of the compounds of iron and other transition elements are due to the presence of a smaller number of electrons in the 3td subshell than required to fill this subshell. For example, ferric ion can have all five of its 2>d electrons with spins oriented in the same direction, because there are five 2>d orbitals in the 3d subshell, and the Pauli principle permits parallel orientation of the spins of electrons so long as there is only one electron per orbital. The ferrous ion. is easily oxidized to ferric ion by air or other oxidizing agents. Both bipositive and terpositive iron form complexes, such as the ferrocyanide ion, Fe(CN)e and the ferricyanide ion, Fe(CN)e, but they do not form complexes with ammonia. 

Inside the tin can Sn" "" " forms complex ions with the organic liquids food. This causes reversal of polarity and tin becomes active and corrodes as Sn++. The corrosion of the base metal (carbon steel) is thus prevented by the sacrificial action of tin coating. The tin ions also inhibit the corrosion of base carbon steel by plating on the iron. Tin is plated on steel as Sn-Fe alloy which has a more noble potential than steel. Thus, within the food container the corrosion of steel is prevented. 

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. 

Thiocyanates are rather stable to air, oxidation, and dilute nitric acid. Of considerable practical importance are the reactions of thiocyanate with metal cations. Silver, mercury, lead, and cuprous thiocyanates precipitate. Many metals form complexes. The deep red complex of ferric iron with thiocyanate, [Fe(SCN)g] , is an effective iadicator for either ion. Various metal thiocyanate complexes with transition metals can be extracted iato organic solvents. 

The thermodynamic data pertinent to the corrosion of metals in aqueous media have been systematically assembled in a form that has become known as Pourbaix diagrams (11). The data include the potential and pH dependence of metal, metal oxide, and metal hydroxide reactions and, in some cases, complex ions. The potential and pH dependence of the hydrogen and oxygen reactions are also suppHed because these are the common corrosion cathodic reactions. The Pourbaix diagram for the iron—water system is given as Figure 1. 

Ferrous ion, iron(II), forms a complex with six cyanide ions, CN- the octahedral complex is called ferrocyanide. Ferric ion, iron(III), forms a complex with six cyanide ions the octahedral complex is called ferricyanide. Write the structural formulas for the ferrocyanide and the ferricyanide complex ions. 

Sulphuric acid is not recommended, because sulphate ions have a certain tendency to form complexes with iron(III) ions. Silver, copper, nickel, cobalt, titanium, uranium, molybdenum, mercury (>lgL-1), zinc, cadmium, and bismuth interfere. Mercury(I) and tin(II) salts, if present, should be converted into the mercury(II) and tin(IV) salts, otherwise the colour is destroyed. Phosphates, arsenates, fluorides, oxalates, and tartrates interfere, since they form fairly stable complexes with iron(III) ions the influence of phosphates and arsenates is reduced by the presence of a comparatively high concentration of acid. 

A stopped flow technique coupled with spectrophotometric analysis of the iron (II) complex formed has been used to investigate - the reactions of some organic complexes of iron(III) with the ion Fe ". The iron(III) was complexed with 1,10-phenanthroline, various substituted 1,10-phenanthrolines (5-methyl-, 5-nitro-, 5-chloro-, 5-phenyl-, 5,6-dimethyl-, 4,7-diphenyl-, and 3,4,7,8-tetramethyl-) and 2,2 -dipyridine, 4,4 -dimethyl-2,2 -dipyridine, and 2,2, 2"-tripyridine. The wavelengths used for the analysis lay in the region 500-552 m/i. 

There is a discrepancy between the cyanide criteria for both aquatic and drinking water standards and the current analytical technology. The criteria are stated for free cyanide (which Includes hydrocyanic acid and the cyanide ion), but the EPA approved analytical methodology for total cyanide measures the free and combined forms (11). This test probably overestimates the potential toxicity. An alternative method (cyanides amenable to chlorination) measures those cyanide complexes which are readily dissociated, but does not measure the iron cyanide complexes which dissociate in sunlight. This method probably tends to underestimate the potential toxicity. Other methods have been proposed, but similar problems exist (12). The Department of Ecology used the EPA-approved APHA procedure which includes a distillation step for the quantification of total cyanide (13,14). A modification of the procedure which omits the distillation step was used for estimation of free cyanide. Later in the study, the Company used a microdiffusion method for free cyanide (15). 

Iron(II) complexes are often included in studies when complexes are prepared from a large number of different metal ions. 2-formylpyridine thiosemicarbazone, 5, forms brown [Fe(5)2A2] (A = Cl, Br) when prepared in ethanol and [Fe(5-H)2] from aqueous alcohol solution [156], All of these complexes are diamagnetic. The resonance Raman and infrared spectra of [Fe(5-H)2] were examined in detail [130] and coordination occurs via the pyridyl nitrogen, azomethine nitrogen and thiol sulfur. There is appreciable d-d sulfur-to-iron(II) Jt-bonding. Solution studies of iron(II) complexes of some 5-substituted-2-formylpyridine thiosemicarbazones have been reported [157], but no solids... 

Bond length differences between HS and LS isomers have been determined for a number of iron(II), iron(III) and cobalt(II) complexes on the basis of multiple temperature X-ray diffraction structure studies [6]. The available results have been collected in Table 17. Average values for the bond length changes characteristic for a particular transition-metal ion have been extracted from these data and are obtained as AR 0.17 A for iron(II) complexes, AR 0.13 A for iron(III) complexes, and AR = 0.06 A for cobalt(II) complexes. These values may be compared with the differences of ionic radii between the HS and LS forms of iron(II), iron(III) and cobalt(II) which were estimated some time ago [184] as 0.16, 0.095, and 0.085 A, respectively. 

---