Chromate and dichromate

Chromate and Dichromate

Properties

Molecular

formula

CrO2−4

and Cr2O2−7

Molar mass 115.994 g mol-1 and 215.988 g

mol-1

Except where noted otherwise, data are given for

materials in their standard state (at 25 °C, 100 kPa)

Infobox references

A sample of potassium chromate

Chromate and dichromateFrom Wikipedia, the free encyclopedia

Chromate salts contain the chromate anion, CrO42−.

Dichromate salts contain the dichromate anion, Cr2O72−.

They are oxoanions of chromium in the oxidation state +6.They are moderately strong oxidizing agents.

Contents

1 Chemical properties

1.1 Acid-base properties

1.2 Oxidation-reduction properties2 Applications

3 Natural occurrence and production4 Safety

5 See also

6 Notes

7 References

8 External links

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 complexChromium(VI) oxide peroxide, CrO(O2)2, is formed; it is

an uncharged covalent molecule which may be extractedinto 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 chemicalequilibrium.

2 CrO42− + 2 H+ Cr2O7

2− + H2O

The predominance diagram shows that the position of the equilibriumdepends on both pH and the analytical concentration of

chromium.[notes 1] The chromate ion is the predominant species inalkaline solutions, but dichromate can become the predominant ion in acidic solutions. The change in colour withpH from yellow (chromate) to orange (dichromate) and the reversible nature of the equilibrium have beenbeautifully illustrated (http://www.youtube.com/watch?v=zP9qEiaL4kQ)

Systematic name

A sample of potassium dichromate

chromate ion

dichromate ion

Predominance diagram for

chromate

Further condensation reactions can occur in strongly acidic solution with the formation of trichromates,

Cr3O102−, and tetrachromates, Cr4O13

2−. All polyoxyanions of chromium(VI) have structures made up of

tetrahedral CrO4 units sharing corners.[2]

The chromate ion is a weak acid.

HCrO4− CrO4

2− + H+; pKa = ca. 5.9

The hydrogen chromate ion, HCrO4-, is also in

equilibrium with the dichromate ion.

2HCrO4− Cr2O7

2− + H2O

This equilibrium does not involve a change in hydrogen ion concentration,so should be independent of pH. The red line on the predominancediagram is not quite horizontal due to the simultaneous equilibrium withthe chromate ion. The hydrogenchromate ion may be protonated, withthe 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

oxidizing 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 fromcorrosion and to improve paint adhesion. Chromate and dichromate salts of heavy metals, lanthanides andalkaline earth metals are only very slightly soluble in water and are thus used as pigments. The lead containing

School bus painted in Chrome

yellow[6]

Crocoite specimen from the

Red Lead Mine, Tasmania,

Australia

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 foundin the Atacama desert.

Chromite ore is heated with a mixture of calcium carbonate andsodium carbonate in the presence of air. The chromium is oxidized to thehexavalent 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 andleaves the insoluble iron oxide. Normally the chromate solution is furtherprocessed 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 (andby market commonality the rest of the world) by EU Parliament directive 2002/95/EC

See also

Chromate conversion coating

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.). Butterworth-Heinemann. p. 637. ISBN 0080379419.

2. ^ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 1009. ISBN 0080379419.

3. ^ IUPAC SC-Database (http://www.acadsoft.co.uk/scdbase/scdbase.htm) A comprehensive database ofpublished 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 (http://dx.doi.org/10.1016%2FS0277-

5387%2897%2900128-9).

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 ofChemicals in the Environment (http://books.google.de/books?id=CjWQ6_7AnI4C&pg=PA13). BNA Books.p. 13. ISBN 978-0-87179-752-0.

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 ofIndustrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a07_067(http://dx.doi.org/10.1002%2F14356007.a07_067).

8. ^ Papp, John F.; Lipin Bruce R. (2006). "Chromite" (http://books.google.de/books?id=zNicdkuulE4C&pg=PA309). Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.).SME. ISBN 978-0-87335-233-8.

External links

National Pollutant Inventory - Chromium VI and compounds fact sheet

(http://www.npi.gov.au/substances/chromium-vi/index.html)

Retrieved from "http://en.wikipedia.org/w/index.php?title=Chromate_and_dichromate&oldid=585148366"

Categories: Chromates Oxidizing agents Oxoanions

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