|Jmol-3D images||Image 1|
|Molar mass||329.24 g/mol|
|Appearance||deep red crystals, sometimes small pellets, orange to dark red powder|
|Density||1.89 g/cm3, solid|
|Melting point||300 °C; 572 °F; 573 K|
|Solubility in water||330 g/L ("cold water")
464 g/L (20°C)
775 g/L ("hot water")
|Solubility||slightly soluble in alcohol
soluble in acid
soluble in water
|octahedral at Fe|
|EU Index||Not listed|
|R-phrases||R20, R21, R22, R32|
|Other anions||Potassium ferrocyanide|
|Other cations||Prussian blue|
| (what is: / ?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C or 77 °F, 100 kPa)
Potassium ferricyanide is the chemical compound with the formula K3[Fe(CN)6]. This bright red salt contains the octahedrally coordinated [Fe(CN)6]3− ion. It is soluble in water and its solution shows some green-yellow fluorescence.
- 2 K4[Fe(CN)6] + Cl2 → 2 K3[Fe(CN)6] + 2 KCl
The compound has widespread use in blueprint drawing and in photography (Cyanotype process). Iron and copper toning involve the use of potassium ferricyanide. Potassium ferricyanide is used as an oxidizing agent to remove silver from negatives and positives, a process called dot etching. In color photography, potassium ferricyanide is used to reduce the size of color dots without reducing their number, as a kind of manual color correction. The compound is also used to harden iron and steel, in electroplating, dyeing wool, as a laboratory reagent, and as a mild oxidizing agent in organic chemistry.
It is also used in black-and-white photography with sodium thiosulfate (hypo) to reduce the density of a negative or gelatin silver print where the mixture is known as Farmer's reducer; this can help offset problems from overexposure of the negative, or brighten the highlights in the print.
Potassium ferricyanide is also one of two compounds present in ferroxyl indicator solution (along with phenolphthalein) which turns blue (Prussian blue) in the presence of Fe2+ ions, and which can therefore be used to detect metal oxidation that will lead to rust. It is possible to calculate the number of moles of Fe2+ ions by using a colorimeter, because of the very intense color of Prussian blue Fe4[Fe(CN)6]3.
Potassium ferricyanide is often used in physiology experiments as a means of increasing a solution's redox potential (E°' ~ 436 mV at pH 7). As such, it can oxidize reduced cytochrome c (E°' ~ 247 mV at pH 7) in intact isolated mitochondria.
Sodium dithionite is usually used as a reducing chemical in such experiments (E°' ~ −420 mV at pH 7).
Potassium ferricyanide is used in many amperometric biosensors as an electron transfer agent replacing an enzyme's natural electron transfer agent such as oxygen as with the enzyme glucose oxidase. It is used as this ingredient in many commercially available blood glucose meters for use by diabetics.
Potassium ferricyanide is combined with potassium hydroxide (or sodium hydroxide as a substitute) and water to formulate Murakami's etchant. This etchant is used by metallographers to provide contrast between binder and carbide phases in cemented carbides.
In histology, potassium ferricyanide is used to detect ferrous iron in biological tissue, with the stain Tirmann Schmeltzer's Turnbull's blue. In this reaction, potassium ferricyanide reacts with ferrous iron in acidic solution to produce an insoluble blue pigment, commonly referred to as Turnbull's blue. To detect ferric (Fe3+) iron, potassium ferrocyanide is used instead in the Perls' Prussian blue staining method. The pigment produced is commonly known as Prussian blue. It has been found that the compound formed in the Turnbull's blue reaction and the compound formed in the Prussian blue reaction are the same unique compound, Prussian blue.
Potassium ferricyanide has very low toxicity, its main hazard being that it is a mild irritant to the eyes and skin. However, under very strongly acidic conditions, highly toxic hydrogen cyanide gas is evolved, according to the equation:
- 6 H+ + [Fe(CN)6]3− → 6 HCN + Fe3+
The reaction with hydrochloric acid is as follows:
- 6 HCl + K3[Fe(CN)6] → 6 HCN + FeCl3 + 3 KCl
- Kwong, H.-L. (2004). "Potassium Ferricyanide". In Paquette, L. Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289.
- Sharpe, A. G. (1976). The Chemistry of Cyano Complexes of the Transition Metals. London: Academic Press.
- "Chemicals". Encyclopedia of Textual Criticism. SkyPoint.
- Stroebel, L.; Zakia, R. D. (1993). "Farmer's Reducer". The Focal Encyclopedia of Photography. Focal Press. p. 297. ISBN 978-0-240-51417-8.
- Dunbar, K. R.; Heintz, R. A. (1997). "Chemistry of Transition Metal Cyanide Compounds: Modern Perspectives". Progress in Inorganic Chemistry 45. pp. 283–391. doi:10.1002/9780470166468.ch4.
- Carson, F. L. (1997). Histotechnology: A Self-Instructional Text (2nd ed.). Chicago: American Society of Clinical Pathologists. pp. 209–211. ISBN 0-89189-411-X.
- Tafesse, F. (2003). "Comparative Studies on Prussian Blue or Diaquatetraamine-Cobalt(III) Promoted Hydrolysis of 4-Nitrophenylphosphate in Microemulsions" (pdf). International Journal of Molecular Sciences 4 (6): 362–370. doi:10.3390/i4060362.
- Verdaguer, M.; Galvez, N.; Garde, R.; Desplanches, C. (2002). "Electrons at Work in Prussian Blue Analogues" (pdf). Electrochemical Society Interface 11 (3): 28–32. doi:10.1002/chin.200304218.
- "MSDS for potassium ferricyanide".