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'[[Image:Potassium-chromate-sample.jpg|thumb|200px|A sample of [[potassium chromate]]]] [[Image:Potassium-dichromate-sample.jpg|thumb|200px|A sample of [[potassium dichromate]]]] '''Chromates''' and '''dichromates''' are [[salt]]s of [[chromic acid]] and dichromic acid, respectively. Chromate salts contain the chromate anion, [[Chromium|Cr]][[Oxygen|O]]{{su|b=4|p=2−}}, and usually have an intense yellow color. Dichromate salts contain the dichromate anion, Cr₂O{{su|b=7|p=2−}}, and usually have an intense orange color. The chromates and dichromates of heavy metals (such as silver and lead) often have a red color. == Characteristics == The [[chromium]] atoms in both of these anions are in [[oxidation state]] +6. The chromate and dichromate ions are fairly strong [[oxidizing agent]]s. In an [[aqueous]] [[Solution (chemistry)|solution]], chromate and dichromate anions exist in a [[chemical equilibrium]]. :2 CrO{{su|b=4|p=2−}} + 2 H₃O⁺ ⇌ Cr₂O{{su|b=7|p=2−}} + 3 H₂O This equilibrium can be pushed toward dichromate by lowering the [[pH]] (making the solution more [[acidic]]) or in the other direction towards chromate by raising the pH to [[base (chemistry)|basic]]. This is a classic example of [[Le Chatelier's principle]] at work. This equilibrium is also dependent on concentration of chromium in solution. At lower pH further condensation to more complex chromates is possible. The chromate and dichromate are strong oxidizing reagents at low pH:<ref name="HollemanAF">{{cite book | publisher = Walter de Gruyter | year = 1985 | edition = 91–100 | pages = 1081–1095 | isbn = 3110075113 | title = Lehrbuch der Anorganischen Chemie | first = Arnold F. | last = Holleman | coauthors = Wiberg, Egon; Wiberg, Nils; | chapter = Chromium| language = German}}</ref> :Cr₂O{{su|b=7|p=2−}} + 14 H₃O⁺ + 6 e⁻ → 2 Cr³⁺ + 21 H₂O (ε₀ = 1.33 V) but only moderate ones at high pH:<ref name="HollemanAF"/> :{{chem|CrO|4|2-}} + 4 {{chem|H|2|O}} + 3 {{chem|e|-}} → {{chem|Cr(OH)|3}} + 5 {{chem|OH|-}} (ε₀ = −0.13 V) ==Applications== Chromate and dichromate salts are widely used in chemical industry, most significantly in the isolation of chromium from its ores.<ref name=Ullmann>Gerd Anger, Jost Halstenberg, Klaus Hochgeschwender, Christoph Scherhag, Ulrich Korallus, Herbert Knopf, Peter Schmidt, Manfred Ohlinger, "Chromium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005.</ref> [[File:Laidlaw school bus.jpg|thumb| right| School bus painted in [[Chrome yellow]]<ref name="Yellow">{{cite book | title = Toxic Substances Controls Guide: Federal Regulation of Chemicals in the Environment | first = Mary Devine | last = Worobec | coauthor = Hogue, Cheryl| page = 13 | BNA Books | year = 1992 | isbn = 9780871797520 | url =http://books.google.de/books?id=CjWQ6_7AnI4C&pg=PA13}}</ref>]] The [[sodium dichromate]] (Na₂Cr₂O₇) and [[potassium dichromate]] (K₂Cr₂O₇), as well as other derivatives are water soluble granular solids and are common [[reagent]]s. Chromate and dichromate salts of [[heavy metals]], [[lanthanides]] and [[alkaline earth metals]] are only very slightly soluble in water and are thus used as pigments. The lead containing pigment [[Chrome Yellow]] was used for a very long time before environmental regulations discouraged its use.<ref name="Yellow"/> Chromates and dichromates are used in [[Chromate conversion coating|chrome plating]] to protect metals for corrosion protection and to improve paint adhesion. When used as oxidizing agents or titrants in a [[redox]] [[chemical reaction]], chromates and dichromates convert into trivalent chromium, Cr³⁺, salts of which typically have a distinctively different blue-green color.<ref name=Ullmann/> ==Production== Chromates are an intermediate in the production of chromium metal. In the production of chromium from chromite ore (FeCr₂O₄) the iron has to be separated from the chromium in a two step roasting and leaching process. The chromite ore is heated with a mixture of [[calcium carbonate]] and [[sodium carbonate]] in the presence of air. The chromium is oxidized to the hexavalent form, while the iron forms the stable Fe₂O₃. The subsequent leaching at higher temperatures dissolves the [[chromate]]s and leaves the unsoluble iron oxide. The chromate is converted by sulfuric acid into the dichromate.<ref name="IndMin">{{cite book | title =Industrial Minerals & Rocks: Commodities, Markets, and Uses | edition = 7th | publisher = SME | year = 2006 | isbn = 9780873352338 | chapter = Chromite | first = John F. | last = Papp | coauthor = Lipin Bruce R. | url = http://books.google.de/books?id=zNicdkuulE4C&pg=PA309 }} </ref> :4 FeCr₂O₄ + 8 Na₂CO₃ + 7 O₂ → 8 Na₂CrO₄ + 2 Fe₂O₃ + 8 CO₂ :Na₂CrO₄ + H₂SO₄ → Na₂Cr₂O₄ + Na₂SO₄ + H₂O To get chromium metal the dichromate is converted to the chromium(III) oxide by reduction with carbon and then reduced in an aluminothermic reaction to chromium.<ref name="IndMin"/> ==Safety== Chromium in the +6 (or VI) oxidation state is often referred to as [[hexavalent chromium]]. Such compounds, especially when air-borne, are [[carcinogenic]]. All hexavalent chromium compounds are considered toxic. The use of chromate compounds in manufactured goods is restricted in the EU (and by market commonality the rest of the world) by EU Parliament directive 2002/95/EC == Structures == {| class="wikitable" style="margin: 1em auto 1em auto" |align="center" | [[Image:Chromate-3D-balls.png|120px]] || align="center" | [[Image:Chromate-2D-dimensions.png|150px]] || align="center" | [[Image:Dichromate-3D-balls.png|150px]] || align="center" | [[Image:Dichromate-2D-dimensions.png|150px]] |- |colspan="2" align="center" |The tetrahedral chromate ion, CrO₄²⁻ || colspan="2" align="center" | The dichromate ion, Cr₂O₇²⁻, consists of two corner-sharing tetrahedra |} == Natural occurrence == [[Image:Crocoite from Tasmania.jpg|thumb|250px|Crocoite specimen from the Red Lead Mine, [[Tasmania]], [[Australia]]]] Chromate minerals are rarely found in the nature. The most commonly met is [[crocoite]]. Potassium-bearing chromates and related compounds are known from the [[Atacama desert]], but are very rare minerals. ==Examples== *[[Ammonium dichromate]] (NH₄)₂Cr₂O₇ *[[Barium chromate]] BaCrO₄ *[[Cadmium chromate]] CdCrO₄ *[[Caesium chromate]] Cs₂CrO₄ *[[Calcium dichromate]] CaCr₂O₇ *[[Cobalt(II) chromate]] CoCrO₄ *[[Cobalt(II) dichromate]] CoCr₂O₇ *[[Cobalt(III) chromate]] Co₂(CrO₄)₃ *[[Cobalt(III) dichromate]] Co₂(Cr₂O₇)₃ *[[Copper(I) chromate]] Cu₂CrO₄ *[[Copper(I) dichromate]] Cu₂Cr₂O₇ *[[Copper(II) chromate]] CuCrO₄ *[[Copper(II) dichromate]] CuCr₂O₇ *[[Iron(III) chromate]] Fe₂(CrO₄)₃ *[[Iron(III) dichromate]] Fe₂(Cr₂O₇)₃ *[[Lead(II) chromate]] PbCrO₄ *[[Lead(II) dichromate]] PbCr₂O₇ *[[Lead(IV) chromate]] Pb(CrO₄)₂ *[[Lead(IV) dichromate]] Pb(Cr₂O₇)₂ *[[Nickel(II) chromate]] NiCrO₄ *[[Nickel(II) dichromate]] NiCr₂O₇ *[[Nickel(III) chromate]] Ni₂(CrO₄)₃ *[[Nickel(III) dichromate]] Ni₂(Cr₂O₇)₃ *[[Potassium chromate]] K₂CrO₄ *[[Potassium trioxochlorochromate]] KCrO₃Cl *[[Potassium dichromate]] K₂Cr₂O₇ *[[Pyridinium chlorochromate]] *[[Silver chromate]] Ag₂CrO₄ *[[Silver dichromate]] Ag₂Cr₂O₇ *[[Sodium chromate]] Na₂CrO₄ *[[Strontium chromate]] SrCrO₄ *[[Tin(II) chromate]] SnCrO₄ *[[Tin(II) dichromate]] SnCr₂O₇ *[[Tin(IV) chromate]] Sn(CrO₄)₂ *[[Tin(IV) dichromate]] Sn(Cr₂O₇)₂ *[[Zinc chromate]] ZnCrO₄ *[[Zinc dichromate]] ZnCr₂O₇ == See also == *[[Chromate conversion coating]] ==References== {{reflist}} == External links == *[http://monographs.iarc.fr/ENG/Monographs/vol49/volume49.pdf IARC Monograph "Chromium and Chromium compounds"] *[http://www.npi.gov.au/database/substance-info/profiles/25.html National Pollutant Inventory - Chromium VI and compounds fact sheet] [[Category:Chromates| ]] [[Category:IARC Group 1 carcinogens]] [[Category:Oxoanions]] [[de:Chromate]] [[es:Cromato]] [[fr:Dichromate]] [[he:כרומט]] [[it:Cromato]] [[ja:クロム酸塩]] [[nl:Chromateren]] [[pt:Cromato]] [[ru:Хроматы]] [[fi:Kromaatti]] [[zh:铬酸盐]]'