Chromate and dichromate

Chromate salts contain the chromate anion, CrO42−. Dichromate salts contain the dichromate anion, Cr2O72−. They are oxyanions of chromium in the oxidation state +6. They are moderately strong oxidizing agents.

Contents

Chemical properties

Chromates react with hydrogen peroxide giving products in which peroxide, O22−, replaces one or more oxygen atoms. In acid solution the unstable blue peroxo complex Chromium(VI) oxide peroxide, CrO(O2)2, is formed; it is an uncharged covalent molecule which may be extracted into ether. Addition of pyridine, results in the formation of the more stable complex CrO(O2)2py.[1]

Acid-base properties

In aqueous solution, chromate and dichromate anions exist in a chemical equilibrium.

2 CrO42− + 2 H+ Cr2O72− + H2O

The predominance diagram shows that the position of the equilibrium depends on both pH and the analytical concentration of chromium.[notes 1] The chromate ion is the predominant species in alkaline solutions, but dichromate can become the predominant ion in acidic solutions. The change in colour with pH from yellow (chromate) to orange (dichromate) and the reversible nature of the equilibrium have been beautifully illustrated

Further condensation reactions can occur in strongly acidic solution with the formation of trichromates, Cr3O102−, and tetrachromates, Cr4O132−. All polyoxyanions of chromium(VI) have structures made up of tetrahedral CrO4 units sharing corners.[2]

The chromate ion is a weak base.

HCrO4− CrO42− + H+; pKa = ca. 5.9

The hydrogenchromate ion, HCrO4-, is also in equilibrium with the dichromate ion.

2HCrO4− Cr2O72− + H2O

This equilibrium does not involve a change in hydrogen ion concentration, so should be independent of pH. The red line on the predominance diagram is not quite horizontal due to the simultaneous equilibrium with the chromate ion. The hydrogenchromate ion may be protonated, with the formation of molecular chromic acid, H2CrO4,but the pKa for the equilibrium

H2CrO4 [HCrO4]− + H+

is not well characterized. Reported values vary between about -0.8 to 1.6.[3]

The dichromate ion is a somewhat weaker base than the chromate ion.

[HCr2O7]− [Cr2O7]2− + H+, pK = 1.8[4]

The pK value for this reaction shows that is can be ignored at pH > 4.

Oxidation-reduction properties

The chromate and dichromate ions are fairly strong oxidizing agents. Commonly three electrons are added to a chromium atom, reducing it to oxidation state +3. In acid solution the aquated Cr3+ ion is produced.

Cr2O72− + 14 H3O+ + 6 e− → 2 Cr3+ + 21 H2O (ε0 = 1.33 V)

In alkaline solution chromium(III) hydroxide is produced. The redox potential shows that chromates are weaker oxidising agent in alkaline solution than in acid solution.[5]

CrO42- + 4 H2O + 3 e− → Cr(OH)3 + 5 OH− (ε0 = −0.13 V)

Applications

Approximately 136,000,000 kilograms (300,000,000 lb) of hexavalent chromium, mainly sodium dichromate, were produced in 1985.[7] Chromates and dichromates are used in chrome plating to protect metals for corrosion protection and to improve paint adhesion. 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.[6] 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.[7]

Natural occurrence and production

The primary chromium ore is the mixed metal oxide chromite, FeCr2O4, found as brittle metallic black crystals or granules. The rare mineral crocoite, PbCrO4, occurs as spectacular long red crystals. Rare potassium chromate minerals and related compounds are found in the Atacama desert.

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 iron(III) oxide, Fe2O3.

4 FeCr2O4 + 8 Na2CO3 + 7 O2 → 8 Na2CrO4 + 2Fe2O3 + 8 CO2

The subsequent leaching at higher temperatures dissolves the chromates and leaves the insoluble iron oxide. Normally the chromate solution is further processed to make chromium metal, but a chromate salt may be obtained directly from the liquor.[8]

Safety

All hexavalent chromium compounds are toxic due to their oxidizing power. They may be carcinogenic, especially when air-borne. 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

See also

Notes

  1. ^ pCr is equal to minus the logarithm of the analytical concentration of chromium. Thus, when pCr=2, the chromium concentration is 10-2 mol dm-3

References

  1. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 637. ISBN 0080379419. 
  2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Oxford: Butterworth-Heinemann. p. 1009. ISBN 0080379419. 
  3. ^ IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands
  4. ^ Brito, F.; Ascanioa, J.; Mateoa, S.; Hernándeza, C.; Araujoa, L.; Gili, P.; Martín-Zarzab, P.; Domínguez, S.; Mederos, A. (1997). "Equilibria of chromate(VI) species in acid medium and ab initio studies of these species". Polyhedron 16 (21): 3835–3846. doi:10.1016/S0277-5387(97)00128-9. 
  5. ^ Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5 
  6. ^ a b Worobec, Mary Devine; Hogue, Cheryl (1992). Toxic Substances Controls Guide: Federal Regulation of Chemicals in the Environment. BNA Books. p. 13. ISBN 9780871797520. http://books.google.de/books?id=CjWQ6_7AnI4C&pg=PA13. 
  7. ^ a b Anger, Gerd; Halstenberg, Jost; Hochgeschwender, Klaus; Scherhag, Christoph, Korallus, Ulrich; Knopf, Herbert; Schmidt, Peter; Ohlinger, Manfred. (2005). "Chromium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. ISBN 10.1002/14356007.a07_067. 
  8. ^ Papp, John F.; Lipin Bruce R. (2006). "Chromite". Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.). SME. ISBN 9780873352338. http://books.google.de/books?id=zNicdkuulE4C&pg=PA309. 

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