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Potassium ferricyanide
CAS number 13746-66-2 Yes check.svgY
PubChem 26250
RTECS number LJ8225000
Molecular formula C6N6FeK3
Molar mass 329.24 g/mol
Appearance deep red crystals
Density 1.89 g/cm3, solid
Melting point

300 °C, 573 K, 572 °F

Boiling point


Solubility in water 33 g/100 mL ("cold water")
46.4 g/100mL (20°C)
77.5 g/100 mL ("hot water")[1]
Solubility slightly soluble in alcohol
soluble in acid
soluble in water
Crystal structure monoclinic
octahedral at Fe
EU Index Not listed
Flash point Non-flammable
Related compounds
Other anions Potassium ferrocyanide
Other cations Prussian blue
 Yes check.svgY (what is this?)  (verify)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Potassium ferricyanide is the chemical compound with the formula K3[Fe(CN)6]. This bright red salt consists of the coordination compound [Fe(CN)6]3−.[2] It is soluble in water and its solution shows some green-yellow fluorescence.

3 K+ HexacyanidoferratIII.svg



Potassium ferricyanide is manufactured by passing chlorine through a solution of potassium ferrocyanide. Potassium ferricyanide separates from the solution:

2 K4[Fe(CN)6] + Cl2 → 2 K3[Fe(CN)6] + 2 KCl


In 19th century, it was used for reading palimpsests and old manuscripts[3].

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 photography with sodium thiosulfate (hypo)[4] to reduce the density of a negative where the mixture is known as Farmer's reducer; this can help offset problems from overexposure. Variants of Farmer's reducer can also be used as the intermediate step in reversal photography to dissolve the silver image produced by the first development.

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 (Eo' ~ 436 mV at pH 7). Sodium dithionite is usually used as a reducing chemical in such experiments (Eo' ~ −420 mV at pH 7).

Potassium ferricyanide is used in many amperometric biosensors as an electron transfer agent replacing an enzymes 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 the main component of Murakami's etchant for cemented carbides.

Prussian blue

Prussian blue, the deep blue pigment in blue printing, is generated by the reaction of K3[Fe(CN)6] with ferrous (Fe2+) ions.[5]

In histology, potassium ferricyanide is used to detect ferrous iron in biological tissue. In this reaction, potassium ferricyanide reacts with ferrous iron in acidic solution to produce an insoluble blue pigment, and both the stain and the pigment are commonly referred to as Turnbull's blue. To detect ferric (Fe3+) iron, potassium ferrocyanide is used instead; the stain and pigment produced are commonly known as Prussian blue.[6] 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.[7][8]


Despite its name, potassium ferricyanide has very low toxicity, its main hazard being that it is a mild irritant to the eyes and skin. However, under acidic conditions, highly toxic hydrogen cyanide gas is evolved, according to the equation:

6H+ + [Fe(CN)6]3− → 6HCN + Fe3+[9]


  1. ^ Kwong, H.-L. "Potassium Ferricyanide" in Encyclopedia of Reagents for Organic Synthesis (Ed: L. Paquette) 2004, J. Wiley & Sons, New York. DOI: 10.1002/047084289.
  2. ^ Sharpe, A. G., The Chemistry of Cyano Complexes of the Transition Metals, Academic Press: London, 1976
  3. ^ Chemicals at the Encyclopedia of Textual Criticism
  4. ^ The Focal Encyclopedia of Photography By Leslie D. Stroebel, Richard D. Zakia [1]
  5. ^ Dunbar, K. R.; Heintz, R. A., "Chemistry of Transition Metal Cyanide Compounds: Modern Perspectives", Progress in Inorganic Chemistry, 1997, volume 45, 283-391.
  6. ^ Carson, Freida L. (1997). Histotechnology: A Self-Instructional Text (2nd ed.), pp. 209-211. Chicago: American Society of Clinical Pathologists. ISBN 0-89189-411-X.
  7. ^ Tafesse, F. (2003). Comparative studies on Prussian blue or diaquatetraamine-cobalt(III) promoted hydrolysis of 4-nitrophenylphosphate in microemulsions. International Journal of Molecular Sciences, 4(6): 362-370.
  8. ^ Verdaguer, M., Galvez, N., Garde, R., & Desplanches, C. (2002). Electrons at work in Prussian blue analogues. Electrochemical Society Interface, 11(3): 28-32.
  9. ^ MSDS for potassium ferricyanide [2]

See also

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