Potassium ferrocyanide
| Names | |
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| IUPAC name
Potassium hexacyanidoferrate(II)
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| Other names | |
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3D model (JSmol)
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| E number | Lua error in Module:Wikidata at line 880: attempt to index field 'wikibase' (a nil value). |
PubChem CID
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| UNII |
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CompTox Dashboard (EPA)
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| Properties | |
| K4[Fe(CN)6] | |
| Molar mass | 368.35 g/mol (anhydrous) 422.388 g/mol (trihydrate) |
| Appearance | Light yellow, crystalline granules |
| Density | 1.85 g/cm3 (trihydrate) |
| Boiling point | (decomposes) |
| trihydrate 28.9 g/100 mL (20 °C) | |
| Solubility | insoluble in ethanol, ether |
| −130.0·10−6 cm3/mol | |
| Hazards | |
| GHS labelling: | |
| GHS09: Environmental hazard | |
| Warning | |
| H411 | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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1.6—3.2g/kg |
| Related compounds | |
Other anions
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Potassium ferricyanide |
Other cations
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Sodium ferrocyanide Prussian blue |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Potassium ferrocyanide is the inorganic compound with formula K4[Fe(CN)6]·3H2O. It is the potassium salt of the coordination complex [Fe(CN)6]4−. This salt forms lemon-yellow monoclinic crystals.
Synthesis
[edit | edit source]In 1752, the French chemist Pierre Joseph Macquer (1718–1784) first reported the preparation of potassium ferrocyanide, which he achieved by reacting Prussian blue (iron(III) ferrocyanide) with potassium hydroxide.[3][4]
Modern production
[edit | edit source]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.[5]
Historical production
[edit | edit source]Historically, the compound was manufactured from nitrogenous organic material, iron filings, and potassium carbonate.[6] 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
[edit | edit source]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.[7]
Upon treatment with chlorine gas, potassium ferrocyanide converts to potassium ferricyanide:
- 2 K4[Fe(CN)6] + Cl2 → 2 K3[Fe(CN)6] + 2 KCl
This reaction can be used to remove potassium ferrocyanide from a solution.[citation needed]
A famous reaction involves treatment with ferric salts, most commonly iron(III) chloride, to give Prussian blue. In the reaction with iron(III) chloride, producing potassium chloride as a side-product:
3 K4[Fe(CN)6] + 4 FeCl3 → Fe4[Fe(CN)6]3 + 12 KCl
With the composition FeIII
4[FeII
(CN)
6]
3, this insoluble but deeply coloured material is the blue of blueprinting, as well as on many famous paintings such as The Great Wave off Kanagawa and The Starry Night.
Applications
[edit | edit source]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.[5]
In the EU, potassium ferrocyanides (E 535–538) were, as of 2017, solely authorised in two food categories as salt additives.
It can also be used in animal feed.[8]
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.[2]
Prior to 1900, before the invention of the Castner process, potassium ferrocyanide was the most important source of alkali metal cyanides.[5] In this historical process, potassium cyanide was produced by decomposing potassium ferrocyanide:[6]
K4[Fe(CN)6] → 4 KCN + FeC2 + N2
Potassium ferrocyanide is a component of white gunpowder or Augendre's powder, a blasting powder of high energy.[9] It comprises 28 parts yellow potassium ferrocyanide, 23 parts cane sugar, and 49 parts of potassium chlorate, mixed under absolute alcohol to prevent premature detonation.[10][11]
Structure
[edit | edit source]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.[12] The K+---NC linkages break when the solid is dissolved in water.[clarification needed][citation needed]
Toxicity
[edit | edit source]The toxicity in rats is low, with lethal dose (LD50) at 1.6—3.2g/kg.[2] The kidneys are the organ for ferrocyanide toxicity.[13]
See also
[edit | edit source]References
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- ^ Lua error in Module:Citation/CS1/Configuration at line 2172: attempt to index field '?' (a nil value). 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.)
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External links
[edit | edit source]- Lua error in Module:Citation/CS1/Configuration at line 2172: attempt to index field '?' (a nil value).
- Lua error in Module:Citation/CS1/Configuration at line 2172: attempt to index field '?' (a nil value).[permanent dead link]
- Photo of large potassium ferrocyanide crystals