Potassium ferrocyanide: Difference between revisions
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{{Short description|Chemical compound}} |
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<b>Potassium ferrocyanide</b> (K<sub>4</sub>Fe(CN)<sub>6</sub>·3H<sub>2</sub>O), also known as yellow prussiate of potash, is a [[complex|coordination compound]] forming lemon-yellow monoclinic crystals at room temperature and decomposing at its boiling point. It is insoluble in alcohol but a litre of water can dissolve just under 300g of the crystals, and the solution can be reduced with acid to release cyanide gas. The resulting [[hydrogen cyanide]] (HCN) boils at 26C and, being lighter than air, quickly evaporates clear of the release point. |
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{{Distinguish|potassium ferricyanide}} |
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{{chembox |
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| Verifiedfields = changed |
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| Watchedfields = changed |
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| verifiedrevid = 444997361 |
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| Name = Potassium hexacyanidoferrate(II) |
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| ImageFile1 = Structure of potassium ferrocyanide.png |
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| ImageFile2 = Potassium Ferrocyanide.png |
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| ImageFile3 = Potassium hexacyanidoferrate(II).jpg |
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| ImageSize3 = 150px |
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| ImageName3 = Potassium ferrocyanide trihydrate |
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| IUPACName = Potassium hexacyanidoferrate(II) |
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| OtherNames = {{Unbulleted list |
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| (Yellow) Prussiate of Potash<ref>{{Cite book|url=https://play.google.com/books/reader?printsec=frontcover&output=reader&id=gXwPAAAAYAAJ&pg=GBS.PA8|title = Five Hundred Useful and Amusing Experiments in Chemistry, and in the Arts and Manufactures: With Observations on the Properties Employed, and Their Application to Useful Purposes|year = 1825|publisher = Thomas Tegg}}</ref> |
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| Potassium hexacyanoferrate (II) trihydrate |
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| Tetrapotassium ferrocyanide trihydrate |
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| Ferrate hexacyano tetrapotassium trihydrate<ref name=JTBaker/> |
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}} |
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| Section1 = {{Chembox Identifiers |
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| CASNo = 13943-58-3 |
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| CASNo_Ref = {{cascite|correct|CAS}} |
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| CASNo_Comment = (anhydrous) |
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| CASNo2_Ref = {{cascite|correct|CAS}} |
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| CASNo2 = 14459-95-1 |
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| CASNo2_Comment = (trihydrate) |
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| ChemSpiderID = 20162028 |
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| EINECS = 237-722-2 |
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| UNII_Ref = {{fdacite|correct|FDA}} |
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| UNII = GTP1P30292 |
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| UNII_Comment = (anhydrous) |
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| UNII2 = 961WP42S65 |
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| UNII2_Comment = (trihydrate) |
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| UNII2_Ref = {{fdacite|correct|FDA}} |
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| PubChem = 161067 |
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| SMILES = [K+].[K+].N#C[Fe-4](C#N)(C#N)(C#N)(C#N)C#N.[K+].[K+] |
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| StdInChI=1S/6CN.Fe.4K.3H2O/c6*1-2;;;;;;;;/h;;;;;;;;;;;3*1H2/q6*-1;+2;4*+1;;; |
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| StdInChIKey = UTYXJYFJPBYDKY-UHFFFAOYSA-N |
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}} |
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| Section2 = {{Chembox Properties |
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| Formula = K<sub>4</sub>[Fe(CN)<sub>6</sub>] |
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| MolarMass = 368.35 g/mol (anhydrous)<br />422.388 g/mol (trihydrate) |
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| Appearance = Light yellow, crystalline granules |
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| Density = 1.85 g/cm<sup>3</sup> (trihydrate) |
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| Solubility = ''trihydrate'' <br /> 28.9 g/100 mL (20 °C) |
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| SolubleOther = insoluble in [[ethanol]], [[diethyl ether|ether]] |
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| MeltingPtC = |
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| BoilingPtC = |
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| BoilingPt_notes = (decomposes) |
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| MagSus = −130.0·10<sup>−6</sup> cm<sup>3</sup>/mol |
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}} |
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| Section7 = {{Chembox Hazards |
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| ExternalSDS = |
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| MainHazards = |
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| FlashPt = Non-flammable |
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| GHSPictograms = {{GHS09}} |
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| GHSSignalWord = Warning |
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| HPhrases = {{H-phrases|411}} |
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| PPhrases = {{P-phrases|}} |
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| NFPA-H = 1 |
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| NFPA-F = 0 |
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| NFPA-R = 0 |
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| NFPA-S = |
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| LD50 = 6400 mg/kg (oral, rat)<ref>https://chem.nlm.nih.gov/chemidplus/rn/13943-58-3 {{Dead link|date=March 2022}}</ref> |
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}} |
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| Section8 = {{Chembox Related |
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| OtherAnions = [[Potassium ferricyanide]] |
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| OtherCations = [[Sodium ferrocyanide]]<br />[[Prussian blue]] |
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}} |
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}} |
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'''Potassium hexacyanidoferrate(II)''' is the [[inorganic compound]] with formula K<sub>4</sub>[Fe(CN)<sub>6</sub>]·3H<sub>2</sub>O. It is the potassium salt of the [[complex (chemistry)|coordination complex]] [Fe(CN)<sub>6</sub>]<sup>4−</sup>. This salt forms lemon-yellow [[monoclinic]] [[crystal]]s. |
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On [[February 20]], [[2002]] four Moroccans were arrested while in possession detailed maps of the US embassy in [[Rome]], the Rome water supply network, and four kilograms of potassium ferrocyanide. |
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==Synthesis== |
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In 1752, the French chemist [[Pierre Macquer|Pierre Joseph Macquer]] (1718–1784) first reported the preparation of Potassium hexacyanidoferrate(II), which he achieved by reacting [[Prussian blue]] (iron(III) ferrocyanide) with [[potassium hydroxide]].<ref>{{cite journal |last1=Macquer |title=Éxamen chymique de bleu de Prusse |journal=Histoire de l'Académie Royale des Sciences …, § Mémoires de l'Académie royale des Sciences |date=1752 |pages=60–77 |url=https://www.biodiversitylibrary.org/item/88112#page/250/mode/1up |trans-title=Chemical examination of Prussian blue |language=fr}} |
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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 [[potassium bitartrate|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.)</ref><ref>{{cite journal |last1=Munroe |first1=Charles E. |last2=Chatard |first2=Thomas M. |title=Manufactures: Chemicals and Allied Products |journal=Twelfth Census of the United States: Bulletins |date=1902 |issue=210 |pages=1–306 |url=https://books.google.com/books?id=8UIUAQAAMAAJ&pg=RA1-PA31}}; see p. 31.</ref> |
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===Modern production=== |
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Potassium hexacyanidoferrate(II) is produced industrially from [[hydrogen cyanide]], [[iron(II) chloride]], and [[calcium hydroxide]], the combination of which affords Ca<sub>2</sub>[Fe(CN)<sub>6</sub>]·11H<sub>2</sub>O. This solution is then treated with potassium salts to precipitate the mixed calcium-potassium salt CaK<sub>2</sub>[Fe(CN)<sub>6</sub>], which in turn is treated with [[potassium carbonate]] to give the tetrapotassium salt.<ref name=Ullmann>{{ cite encyclopedia |author1=Gail, E. |author2=Gos, S. |author3=Kulzer, R. |author4=Lorösch, J. |author5=Rubo, A. |author6=Sauer, M. |author7=Kellens, R. |author8=Reddy, J. |author9=Steier, N. |author10=Hasenpusch, W. | title = Cyano Compounds, Inorganic | encyclopedia = Ullmann's Encyclopedia of Industrial Chemistry |date=October 2011 | publisher = Wiley-VCH | location = Weinheim | doi = 10.1002/14356007.a08_159.pub3 |isbn=978-3527306732 }}</ref> |
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===Historical production=== |
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Historically, the compound was manufactured from nitrogenous organic material, iron filings, and potassium carbonate.<ref name=Wagner>{{cite book|last=Von Wagner|first=Rudolf|title=Manual of chemical technology|year=1897|publisher=D. Appleton & Co.|location=New York|page=474 & 477|url=https://archive.org/details/manualofchemical00wagnuoft}}</ref> Common [[nitrogen]] and [[carbon]] sources were [[torrefaction|torrified]] horn, leather scrap, [[offal]], or dried blood. It was also obtained commercially from gasworks spent oxide (purification of city gas from hydrogen cyanide). |
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==Chemical reactions== |
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Treatment of potassium hexacyanidoferrate(II) with [[nitric acid]] gives H<sub>2</sub>[Fe(NO)(CN)<sub>5</sub>]. After neutralization of this intermediate with [[sodium carbonate]], red crystals of [[sodium nitroprusside]] can be selectively crystallized.<ref>{{cite encyclopedia | author = Seel, F. | title = Sodium nitrosyl cyanoferrate | encyclopedia = Handbook of Preparative Inorganic Chemistry | edition = 2nd | editor = Brauer, G. | publisher = Academic Press | year = 1965 | location = New York | volume = 2 | page = 1768 | lccn = 63-14307 | url = https://www.scribd.com/doc/27443280/Handbook-of-Preparative-Inorganic-Chemistry-Vol-2-2d-Ed-George-Brauer | access-date = 2017-09-10 | archive-date = 2010-03-07 | archive-url = https://web.archive.org/web/20100307011411/http://www.scribd.com/doc/27443280/Handbook-of-Preparative-Inorganic-Chemistry-Vol-2-2d-Ed-George-Brauer | url-status = dead }}</ref> |
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Upon treatment with [[chlorine]] gas, potassium hexacyanidoferrate(II) converts to [[potassium hexacyanidoferrate(III)]]: |
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:2 K<sub>4</sub>[Fe(CN)<sub>6</sub>] + Cl<sub>2</sub> → 2 K<sub>3</sub>[Fe(CN)<sub>6</sub>] + 2 KCl |
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This reaction can be used to remove potassium hexacyanidoferrate(II) from a solution.{{citation needed|date=April 2012}} |
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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: |
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3 K<sub>4</sub>[Fe(CN)<sub>6</sub>] + 4 FeCl<sub>3</sub> → Fe<sub>4</sub>[Fe(CN)<sub>6</sub>]<sub>3</sub> + 12 KCl |
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With the composition Fe{{su|p=III|b=4}}[Fe{{su|p=II}}([[Cyanide|CN]]){{su|b=6}}]{{su|b=3}}, this insoluble but deeply coloured material is the blue of [[blueprint]]ing, as well as on many famous paintings such as [[The Great Wave off Kanagawa]] and [[The Starry Night]]. |
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==Applications== |
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Potassium hexacyanidoferrate(II) finds many niche applications in industry. It and [[Sodium ferrocyanide|the related sodium salt]] are widely used as anticaking agents for both road salt and table salt. The potassium and sodium hexacyanidoferrates(II) are also used in the purification of tin and the separation of copper from molybdenum ores. Potassium hexacyanidoferrate(II) is used in the production of wine and citric acid.<ref name=Ullmann/> |
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In the EU, hexacyanidoferrates(II) (E 535–538) were, as of 2017, solely authorised in two food categories as salt additives. |
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It can also be used in animal feed.<ref>{{cite web |title=EuSalt Expert Meeting on E 535 and E 536 as Feed Additives |url=https://eusalt.com/events/eusalt-expert-meeting-e-535-and-e-536-feed-additives |publisher=EUSalt |access-date=2018-12-06 |archive-date=2019-05-12 |archive-url=https://web.archive.org/web/20190512013808/https://eusalt.com/events/eusalt-expert-meeting-e-535-and-e-536-feed-additives |url-status=dead }}</ref> |
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In the laboratory, potassium hexacyanidoferrate(II) is used to determine the concentration of [[potassium permanganate]], a compound often used in [[titration]]s based on [[redox]] reactions. Potassium hexacyanidoferrate(II) 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. |
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Potassium hexacyanidoferrate(II) can be used as a fertilizer for plants.{{citation needed|date=April 2012}} |
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Prior to 1900, before the invention of the [[Castner process]], potassium hexacyanidoferrate(II) was the most important source of [[alkali metal]] [[Sodium cyanide|cyanides]].<ref name=Ullmann /> In this historical process, [[potassium cyanide]] was produced by decomposing potassium hexacyanidoferrate(II):<ref name=Wagner /> |
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K<sub>4</sub>[Fe(CN)<sub>6</sub>] → 4 KCN + FeC<sub>2</sub> + N<sub>2</sub> |
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==Structure== |
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Like other metal cyanides, solid potassium hexacyanidoferrate(II), both as the hydrate and anhydrous salts, has a complicated polymeric structure. The polymer consists of octahedral [Fe(CN)<sub>6</sub>]<sup>4−</sup> centers crosslinked with K<sup>+</sup> ions that are bound to the CN [[ligand]]s.<ref>{{Cite journal |last1=Willans |first1=Mathew J. |last2=Wasylishen |first2=Roderick E. |last3=McDonald |first3=Robert |date=2009-05-18 |title=Polymorphism of Potassium Ferrocyanide Trihydrate as Studied by Solid-State Multinuclear NMR Spectroscopy and X-ray Diffraction |url=https://pubs.acs.org/doi/10.1021/ic802134j |journal=Inorganic Chemistry |language=en |volume=48 |issue=10 |pages=4342–4353 |doi=10.1021/ic802134j |pmid=19425611 |issn=0020-1669}}</ref> The K<sup>+</sup>---NC linkages break when the solid is dissolved in water.{{clarify|date=May 2020}}{{citation needed|date=May 2020}} |
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===Toxicity=== |
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The toxicity in rats is low, with [[lethal dose]] (LD<sub>50</sub>) at 6400 mg/kg.<ref name=JTBaker>{{cite web | url = http://hazard.com/msds/mf/baker/baker/files/p5763.htm | title = POTASSIUM FERROCYANIDE MSDS Number: P5763 - Effective Date: 12/08/96 | publisher = J. T. Baker Inc. | access-date = 2012-04-08 | url-status = dead | archive-date = 2015-11-21 | archive-url = https://web.archive.org/web/20151121093949/http://hazard.com/msds/mf/baker/baker/files/p5763.htm }}</ref>{{Better source needed|date=July 2024}} The kidneys are the organ for ferrocyanide toxicity.<ref>{{cite journal |last1=Peter Aggett, Fernando Aguilar, Riccardo Crebelli, Birgit Dusemund, Metka Filipič, Maria Jose Frutos, Pierre Galtier, David Gott, Ursula Gundert-Remy, Gunter Georg Kuhnle, Claude Lambré, Jean-Charles Leblanc, Inger Therese Lillegaard, Peter Moldeus, Alicja Mortensen, Agneta Oskarsson, Ivan Stankovic, Ine Waalkens-Berendsen, Rudolf Antonius Woutersen, Matthew Wright and Maged Younes. |title=Re-evaluation of sodium ferrocyanide (E 535), potassium ferrocyanide (E 536) and calcium ferrocyanide (E 538) as food additives |journal=EFSA Journal |date=2018 |volume=16 |issue=7 |page=5374 |doi=10.2903/j.efsa.2018.5374 |pmid=32626000 |pmc=7009536 |url=https://www.efsa.europa.eu/en/efsajournal/pub/5374|doi-access=free }}</ref> |
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==See also== |
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*[[Ferrocyanide]] |
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*[[Potassium hexacyanidoferrate(III)]]<ref>{{Cite journal |last1=Kosugi |first1=Nobuhiro |last2=Yokoyama |first2=Toshihiko |last3=Kuroda |first3=Haruo |date=May 1986 |title=Polarization dependence of XANES of square-planar Ni(CN)2−4 ion. A comparison with octahedral Fe(CN)4−6 and Fe(CN)3−6 ions |url=http://dx.doi.org/10.1016/0301-0104(86)85034-0 |journal=Chemical Physics |volume=104 |issue=3 |pages=449–453 |doi=10.1016/0301-0104(86)85034-0 |issn=0301-0104}}</ref> |
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*[[Ferricyanide]] |
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==References== |
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{{Reflist}} |
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==External links== |
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* {{ cite web | url = http://www.npi.gov.au/substances/cyanide/index.html | publisher = National Pollutant Inventory Australia | title = Cyanide (inorganic) compounds fact sheet }} |
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* {{cite web | url = http://blogs.rediff.com/thinko/2019/07/31/customers-can-rely-on-tata-salt-as-report-says-it-is-safe-to-consume/ | work = rediff.com | title = Potassium Ferrocyanide in Salt Is Entirely Safe To Consume }}{{Dead link|date=October 2022 |bot=InternetArchiveBot |fix-attempted=yes }} |
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{{Potassium compounds}} |
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{{DEFAULTSORT:Potassium Ferrocyanide}} |
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[[Category:Potassium compounds]] |
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[[Category:Iron(II) compounds]] |
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[[Category:Cyano complexes]] |
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[[Category:E-number additives]] |
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[[Category:Nephrotoxins]] |
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[[Category:Iron complexes]] |
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Latest revision as of 02:31, 13 December 2024
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| Names | |
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| IUPAC name
Potassium hexacyanidoferrate(II)
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| Other names | |
| Identifiers | |
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3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.034.279 |
| EC Number |
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| E number | E536 (acidity regulators, ...) |
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: | |
| |
| Warning | |
| H411 | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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6400 mg/kg (oral, rat)[3] |
| 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 hexacyanidoferrate(II) 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]In 1752, the French chemist Pierre Joseph Macquer (1718–1784) first reported the preparation of Potassium hexacyanidoferrate(II), which he achieved by reacting Prussian blue (iron(III) ferrocyanide) with potassium hydroxide.[4][5]
Modern production
[edit]Potassium hexacyanidoferrate(II) 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
[edit]Historically, the compound was manufactured from nitrogenous organic material, 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
[edit]Treatment of potassium hexacyanidoferrate(II) 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]
Upon treatment with chlorine gas, potassium hexacyanidoferrate(II) converts to potassium hexacyanidoferrate(III):
- 2 K4[Fe(CN)6] + Cl2 → 2 K3[Fe(CN)6] + 2 KCl
This reaction can be used to remove potassium hexacyanidoferrate(II) 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]Potassium hexacyanidoferrate(II) 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 hexacyanidoferrates(II) are also used in the purification of tin and the separation of copper from molybdenum ores. Potassium hexacyanidoferrate(II) is used in the production of wine and citric acid.[6]
In the EU, hexacyanidoferrates(II) (E 535–538) were, as of 2017, solely authorised in two food categories as salt additives.
It can also be used in animal feed.[9]
In the laboratory, potassium hexacyanidoferrate(II) is used to determine the concentration of potassium permanganate, a compound often used in titrations based on redox reactions. Potassium hexacyanidoferrate(II) 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 hexacyanidoferrate(II) can be used as a fertilizer for plants.[citation needed]
Prior to 1900, before the invention of the Castner process, potassium hexacyanidoferrate(II) was the most important source of alkali metal cyanides.[6] In this historical process, potassium cyanide was produced by decomposing potassium hexacyanidoferrate(II):[7]
K4[Fe(CN)6] → 4 KCN + FeC2 + N2
Structure
[edit]Like other metal cyanides, solid potassium hexacyanidoferrate(II), 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
[edit]The toxicity in rats is low, with lethal dose (LD50) at 6400 mg/kg.[2][better source needed] The kidneys are the organ for ferrocyanide toxicity.[11]
See also
[edit]References
[edit]- ^ Five Hundred Useful and Amusing Experiments in Chemistry, and in the Arts and Manufactures: With Observations on the Properties Employed, and Their Application to Useful Purposes. Thomas Tegg. 1825.
- ^ a b "POTASSIUM FERROCYANIDE MSDS Number: P5763 - Effective Date: 12/08/96". J. T. Baker Inc. Archived from the original on 2015-11-21. Retrieved 2012-04-08.
- ^ https://chem.nlm.nih.gov/chemidplus/rn/13943-58-3 [dead link]
- ^ 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.)
- ^ Munroe, Charles E.; Chatard, Thomas M. (1902). "Manufactures: Chemicals and Allied Products". Twelfth Census of the United States: Bulletins (210): 1–306.; see p. 31.
- ^ a b c Gail, E.; Gos, S.; Kulzer, R.; Lorösch, J.; Rubo, A.; Sauer, M.; Kellens, R.; Reddy, J.; Steier, N.; Hasenpusch, W. (October 2011). "Cyano Compounds, Inorganic". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a08_159.pub3. ISBN 978-3527306732.
- ^ a b Von Wagner, Rudolf (1897). Manual of chemical technology. New York: D. Appleton & Co. p. 474 & 477.
- ^ 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. LCCN 63-14307. Archived from the original on 2010-03-07. Retrieved 2017-09-10.
- ^ "EuSalt Expert Meeting on E 535 and E 536 as Feed Additives". EUSalt. Archived from the original on 2019-05-12. Retrieved 2018-12-06.
- ^ Willans, Mathew J.; Wasylishen, Roderick E.; McDonald, Robert (2009-05-18). "Polymorphism of Potassium Ferrocyanide Trihydrate as Studied by Solid-State Multinuclear NMR Spectroscopy and X-ray Diffraction". Inorganic Chemistry. 48 (10): 4342–4353. doi:10.1021/ic802134j. ISSN 0020-1669. PMID 19425611.
- ^ Peter Aggett, Fernando Aguilar, Riccardo Crebelli, Birgit Dusemund, Metka Filipič, Maria Jose Frutos, Pierre Galtier, David Gott, Ursula Gundert-Remy, Gunter Georg Kuhnle, Claude Lambré, Jean-Charles Leblanc, Inger Therese Lillegaard, Peter Moldeus, Alicja Mortensen, Agneta Oskarsson, Ivan Stankovic, Ine Waalkens-Berendsen, Rudolf Antonius Woutersen, Matthew Wright and Maged Younes. (2018). "Re-evaluation of sodium ferrocyanide (E 535), potassium ferrocyanide (E 536) and calcium ferrocyanide (E 538) as food additives". EFSA Journal. 16 (7): 5374. doi:10.2903/j.efsa.2018.5374. PMC 7009536. PMID 32626000.
{{cite journal}}: CS1 maint: multiple names: authors list (link) - ^ Kosugi, Nobuhiro; Yokoyama, Toshihiko; Kuroda, Haruo (May 1986). "Polarization dependence of XANES of square-planar Ni(CN)2−4 ion. A comparison with octahedral Fe(CN)4−6 and Fe(CN)3−6 ions". Chemical Physics. 104 (3): 449–453. doi:10.1016/0301-0104(86)85034-0. ISSN 0301-0104.
External links
[edit]- "Cyanide (inorganic) compounds fact sheet". National Pollutant Inventory Australia.
- "Potassium Ferrocyanide in Salt Is Entirely Safe To Consume". rediff.com.[permanent dead link]