Stability of Iron Blue

Iron Blue is an extremely acid-resistant, but base-decomposing pigment.357 Hydrogen cyanide is only released by warm, diluted sulfuric acid, while hydrochloric acid, by contrast, has no effect.358 In a clearly alkaline environment, i.e., in the presence of high concentrations of OH ions, these displace the cyanide ion from the iron(III)-ion. Fe(OH)3 is then precipitated ( rust sludge ), and the Iron Blue is destroyed.359 

359 Iron(III)-hydroxide is even less soluble in this range than Iron Blue on the solubility of Fe(OH)3 see chapter 6.6.3. to be exact, Iron Blue is not totally destroyed at a high pH rather, 

The literature contains authenticated cases of studies with iron at pH values of 9 and 10, in which it is still stable.360 The pH range around 10 to 11 can be considered the critical limit for the stability of Iron Blue. Based on the alkaline behavior of fresh mortar and concrete (in this regard, see also chapter 6.7.2), Iron Blue is only used to paint these surfaces to a limited extent.361 

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Concrete, reliable values on the solubility of Iron Blue are not recorded in the scientific literature. Based on comparative calculations between the known solubility of Fe(OH)3 on the one hand, and the limit value of the pH stability of Iron Blue on the other hand (pH 10), the approximate solubility of Iron Blue in water can, however, be calculated (see chapter 6.6.2.2.). It amounts to approximately 10"24 g Iron Blue per liter of water, this means that 0.000000000000000000000001 g Iron Blue dissolve in 1,000 g of water. 

In the appendix of his expert report under the heading Danksa-gung (Acknowledgements), Rudolf had thanked a number of persons and institutions who had helped him in many ways in the collection of data or sources, the recovery and analysis of samples, or for any assistance in the production of the report. Among these were the firms DEGUSSA AG and Institute Fresenius, since the first had supplied important technical data on the stability of Iron Blue and the second had analyzed most of the samples in Rudolf s presence and initially with his help. Such acknowledgements are usual in scientific publica-... 

Therefore, they did not determine the solubility of Iron Blue, but the measure of stability of the dispersion of fresh precipitations of the pigment. 

It is safe to say that Iron Blue is stable at a pH value of 7, i.e., in a neutral aqueous medium, so we take this as a minimum value. As mentioned earlier, a pH value of about 10 can be considered the upper limit of stability for Iron Blue, so we take this as maximum value for the following calculations. At pH=7, and even more so at pH=10, the free iron concentration is extremely low, since Fe(OH)3 is nearly insoluble (see Table 5). 

Graph 6 Free Fe3+ concentration as a function of pH value and the resulting minimal pKs value of Iron Blue, depending on its stability at the corresponding pH value. pKs value acc. to Tananaev et al. 40.5 according to reflections made here greater than 123, smaller than... 

In the fourth beaker iron(III) salts form deep red complexes such as [Fe(SCN)(H20)5] with the SCN ions. The extreme stability of Prussian blue dominates in the fifth beaker, so that the deep blue color brings the series to a close. Deviations from the given concentrations can lead to slight differences in the effects due to the formation of precipitates or mixed colors. 

Variamine blue (C.I. 37255). The end point in an EDTA titration may sometimes be detected by changes in redox potential, and hence by the use of appropriate redox indicators. An excellent example is variamine blue (4-methoxy-4 -aminodiphenylamine), which may be employed in the complexometric titration of iron(III). When a mixture of iron(II) and (III) is titrated with EDTA the latter disappears first. As soon as an amount of the complexing agent equivalent to the concentration of iron(III) has been added, pFe(III) increases abruptly and consequently there is a sudden decrease in the redox potential (compare Section 2.33) the end point can therefore be detected either potentiometrically or with a redox indicator (10.91). The stability constant of the iron(III) complex FeY- (EDTA = Na2H2Y) is about 1025 and that of the iron(II) complex FeY2 - is 1014 approximate calculations show that the change of redox potential is about 600 millivolts at pH = 2 and that this will be almost independent of the concentration of iron(II) present. The jump in redox potential will also be obtained if no iron(II) salt is actually added, since the extremely minute amount of iron(II) necessary is always present in any pure iron(III) salt. 

The stabilization of Fe by 1,4,7-triathiacyclononane, ([9]-ane-S3, ttcn= (297), has permitted the preparation of the first authentic example of a Turnbull s Blue, i.e., an iron(II)-hexacyano-ferrate(lll) combination, in the form of [Fe(ttcn)2][Fe(CN)6] 2H20. This so-called Ukrainian Red, named in honor of the country of origin of several of the authors, is a valence-trapped (Robin and Day class compound. The redox potential for the [Fe(ttcn)2] " couple is... 

History. Braun and Tschemak [23] obtained phthalocyanine for the first time in 1907 as a byproduct of the preparation of o-cyanobenzamide from phthalimide and acetic anhydride. However, this discovery was of no special interest at the time. In 1927, de Diesbach and von der Weid prepared CuPc in 23 % yield by treating o-dibromobenzene with copper cyanide in pyridine [24], Instead of the colorless dinitriles, they obtained deep blue CuPc and observed the exceptional stability of their product to sulfuric acid, alkalis, and heat. The third observation of a phthalocyanine was made at Scottish Dyes, in 1929 [25], During the preparation of phthalimide from phthalic anhydride and ammonia in an enamel vessel, a greenish blue impurity appeared. Dunsworth and Drescher carried out a preliminary examination of the compound, which was analyzed as an iron complex. It was formed in a chipped region of the enamel with iron from the vessel. Further experiments yielded FePc, CuPc, and NiPc. It was soon realized that these products could be used as pigments or textile colorants. Linstead et al. at the University of London discovered the structure of phthalocyanines and developed improved synthetic methods for several metal phthalocyanines from 1929 to 1934 [1-11]. The important CuPc could not be protected by a patent, because it had been described earlier in the literature [23], Based on Linstead s work the structure of phthalocyanines was confirmed by several physicochemical measurements [26-32], Methods such as X-ray diffraction or electron microscopy verified the planarity of this macrocyclic system. Properties such as polymorphism, absorption spectra, magnetic and catalytic characteristics, oxidation and reduc-... 

In summary, it may be stated that Iron Blue having formed in the interior of a wall as a component of the wall itself, possesses a longevity comparable to the iron oxide from which it has formed. This means simply that Iron Blue possesses a degree of stability which is comparable to that of the masonry itself the Iron Blue will remain contained in the wall for as long as the wall itself remains in existence. 386... 

Chrome greens are coprecipitates of lead chromate yellow and iron blue. As such, they share the limitations in performance of conventional chrome yellows and iron blues, namely alkali sensitivity and limited heat stability. Consequently, they are at best of limited utility in plastics applications. They range in hue from a very light green to a very dark green. 

The. exceptionally high affinity255 of iron(II) for cyanide ion is reflected in the heat of formation256-257 of the [Fe(CN)6]4- anion (equation 31). The stability of this, the ferrocyanide ion, is illustrated by the nature and history258 of its iron(III) salt, Prussian blue, possibly the first isolated coordination complex, which is discussed in Section 44.1.5.2.4. 

In wines, traces of iron, which are picked up, perhaps, from processing and/or storage, or copper, which are picked up from mildew sprays, such as Bordeaux mixture, affect the oxidative stability of wines by acting as the redox shuttles as they transfer between oxidation states. Winemakers discovered that adding ferricyanide to wine, in a process known as blue fining, precipitates copper and iron and thereby reduces their concentrations below 1 ppm, which is considered to be acceptable. Critical control of ferricyanide addition is necessary, as cyanide is also a contaminant that must be measured. Where vineyards have replaced cherry and apple orchards, low concentrations of arsenic have started to appear but they are present at very low concentrations in high quality wines. The arsenic appears from arsenical compounds such as lead and calcium arsenates that were used for many decades as pesticides on apples and cherry orchards. 

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