Potassium ferrocyanide is produced industrially from hydrogen cyanide, iron(II) chloride, and calcium hydroxide, the combination of which affords Ca2[Fe(CN)6]·11H2O. This solution is then treated with potassium salts to precipitate the mixed calcium-potassium salt CaK2[Fe(CN)6], which in turn is treated with potassium carbonate to give the tetrapotassium salt.[6]
Historical production
Historically, the compound was manufactured from organic compounds containing nitrogen, iron filings, and potassium carbonate.[7] Common nitrogen and carbon sources were torrified horn, leather scrap, offal, or dried blood. It was also obtained commercially from gasworks spent oxide (purification of city gas from hydrogen cyanide).
Chemical reactions
Treatment of potassium ferrocyanide with nitric acid gives H2[Fe(NO)(CN)5]. After neutralization of this intermediate with sodium carbonate, red crystals of sodium nitroprusside can be selectively crystallized.[8]
This reaction can be used to remove potassium ferrocyanide from a solution.[citation needed]
A famous reaction involves treatment with ferric salts to give Prussian blue. With the composition FeIII 4[FeII (CN) 6] 3, this insoluble but deeply coloured material is the blue of blueprinting.
Applications
Potassium ferrocyanide finds many niche applications in industry. It and the related sodium salt are widely used as anticaking agents for both road salt and table salt. The potassium and sodium ferrocyanides are also used in the purification of tin and the separation of copper from molybdenum ores. Potassium ferrocyanide is used in the production of wine and citric acid.[6]
In the EU, ferrocyanides (E 535–538) were, as of 2017, solely authorised in two food categories as salt additives.
In the laboratory, potassium ferrocyanide is used to determine the concentration of potassium permanganate, a compound often used in titrations based on redox reactions. Potassium ferrocyanide is used in a mixture with potassium ferricyanide and phosphate buffered solution to provide a buffer for beta-galactosidase, which is used to cleave X-Gal, giving a bright blue visualization where an antibody (or other molecule), conjugated to Beta-gal, has bonded to its target. On reacting with Fe(3) it gives a Prussian blue colour. Thus it is used as an identifying reagent for iron in labs.
Potassium ferrocyanide can be used as a fertilizer for plants.[citation needed]
Prior to 1900, before the invention of the Castner process, potassium ferrocyanide was the most important source of alkali metalcyanides.[6] In this historical process, potassium cyanide was produced by decomposing potassium ferrocyanide:[7]
K4[Fe(CN)6] → 4 KCN + FeC2 + N2
Structure
Like other metal cyanides, solid potassium ferrocyanide, both as the hydrate and anhydrous salts, has a complicated polymeric structure. The polymer consists of octahedral [Fe(CN)6]4− centers crosslinked with K+ ions that are bound to the CN ligands.[10] The K+---NC linkages break when the solid is dissolved in water.[clarification needed][citation needed]
Toxicity
Potassium ferrocyanide is nontoxic, and does not decompose into cyanide in the body. The toxicity in rats is low, with lethal dose (LD50) at 6400 mg/kg.[2] The kidneys are the organ for ferrocyanide toxicity.[11]
^Macquer (1752). "Éxamen chymique de bleu de Prusse" [Chemical examination of Prussian blue]. Histoire de l'Académie Royale des Sciences …, § Mémoires de l'Académie royale des Sciences (in French): 60–77.
From pp. 63-64: "Après avoir essayé ainsi inutilement de décomposer le bleu de Prusse par les acides, … n'avoit plus qu'une couleur jaune un peu rousse." (After having tried so vainly to decompose Prussian blue by acids, I made recourse to alkalies. I put a half ounce of this [Prussian] blue in a flask, and I poured on it ten ounces of a solution of nitre fixed by tartar [i.e., potassium nitrate (nitre) which is mixed with crude cream of tartar and then ignited, producing potassium carbonate]. As soon as these two substances had been mixed together, I saw with astonishment that, without the aid of heat, the blue color had entirely disappeared; the powder [i.e., precipitate] at the bottom of the flask had only a rather gray color: having put this vessel on a sand bath in order to heat the solution until it simmered, this gray color also disappeared entirely, and all that was contained in the flask, both the powder [i.e., precipitate] and the solution, had only a yellow color [that was] a little red.)
^Seel, F. (1965). "Sodium nitrosyl cyanoferrate". In Brauer, G. (ed.). Handbook of Preparative Inorganic Chemistry. Vol. 2 (2nd ed.). New York: Academic Press. p. 1768. LCCN63-14307. Archived from the original on 2010-03-07. Retrieved 2017-09-10.
