Potassium dichromate

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Potassium dichromate
Potassium dichromate
Potassium dichromate
2 K+ [O3Cr−O−CrO3]2−
Names
IUPAC name
Potassium dichromate(VI)
Other names
  • potassium bichromate
  • bichromate of potash
  • dipotassium dichromate
  • dichromic acid, dipotassium salt
  • chromic acid, dipotassium salt
  • lópezite[1]
Identifiers
3D model (JSmol)
ChEMBL
ChemSpider
ECHA InfoCard 100.029.005
EC Number
  • 231-906-6
RTECS number
  • HX7680000
UNII
UN number 3288
  • InChI=1S/2Cr.2K.7O/q;;2*+1;;;;;;2*-1 checkY
    Key: KMUONIBRACKNSN-UHFFFAOYSA-N checkY
  • [K+].[K+].[O-][Cr](=O)(=O)O[Cr]([O-])(=O)=O
Properties[2]
K2Cr2O7
Molar mass 294.182 g·mol−1
Appearance red-orange crystalline solid
Odor odorless
Density 2.68 g/cm3, solid
Melting point 398 °C (748 °F; 671 K)
Boiling point 500 °C (932 °F; 773 K) decomposes
15.1 g/100g
1.738
Structure
Triclinic (α-form, <241.6 °C (466.9 °F; 514.8 K))
Tetrahedral (for Cr)
Thermochemistry
219 J/mol[3]
291.2 J/(K·mol)
−2033 kJ/mol
Hazards
GHS labelling:[4]
GHS03: OxidizingGHS05: CorrosiveGHS06: ToxicGHS07: Exclamation markGHS08: Health hazardGHS09: Environmental hazard
Danger
H272, H301, H312, H314, H317, H330, H334, H335, H340, H350, H360, H372, H410
P201, P202, P210, P220, P221, P260, P264, P270, P271, P272, P273, P280, P284, P301+P310+P330, P301+P330+P331, P303+P361+P353, P304+P340+P310, P305+P351+P338+P310, P308+P313, P333+P313, P342+P311, P363, P370+P378, P391, P403+P233, P405, P501
NFPA 704 (fire diamond)
[5]
NFPA 704 four-colored diamond
4
0
1
OX
0.2 μg/m3 (TWA), 0.5 μg/m3 Skin[5] (STEL)
Lethal dose or concentration (LD, LC):
25 mg/kg (oral, rat)[6]
0.083 mg/L (Rat, female, inhalation, dust/mist)[4]
NIOSH (US health exposure limits):[7]
PEL (Permissible)
 5 μg/m3 (TWA, as CrO3)
REL (Recommended)
 0.2 μg/m3 (TWA, as Cr)
IDLH (Immediate danger)
 15 mg/m3 (as Cr(VI))
Related compounds
Other anions
Other cations
Related compounds
 Potassium permanganate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
checkY verify (what is checkY☒N ?)

Potassium dichromate is the inorganic compound with the formula K2Cr2O7. An orange solid, it is used in diverse laboratory and industrial applications. As with all hexavalent chromium compounds, it is chronically harmful to health. It is a crystalline ionic solid with a very bright, red-orange color. The salt is popular in laboratories because it is not deliquescent, in contrast to the more industrially relevant salt sodium dichromate.[8]

Production

Potassium dichromate is usually prepared by the reaction of sodium dichromate and potassium chloride in the presence of an acid.[8] Potassium dichromate precipitates, because it has a lower solubility than the corresponding sodium salt. Alternatively, it can be also obtained from potassium chromate by roasting chromite ore with potassium hydroxide:

FeCr2O4 + 2 KOH + 1.5 O2 → K2Cr2O7 + Fe(OH)2

Structure

Unit cell of potassium dichromate

X-ray crystallography shows that it is a salt. Two polymorphs are known, but the structure of the dichromate anion is the same in each. It ionizes when dissolved in water, releasing Cr2O2−7. Thus in aqueous solution, it is the dichromate ion that matters in terms of chemical reactions and environmental impact. This anion is a corner-shared bitetrahedron, resembling pyrophosphate.[9]

Reactions

Inorganic

When heated strongly, it decomposes with the evolution of oxygen.

4 K2Cr2O7 → 4 K2CrO4 + 2 Cr2O3 + 3 O2

When an alkali is added to an orange-red solution containing dichromate ions, a yellow solution is obtained due to the formation of chromate ions (CrO2−4). For example, potassium chromate is produced industrially using potassium carbonate:[8]

K2Cr2O7 + K2CO3 → 2 K2CrO4 + CO2

Treatment with cold sulfuric acid gives red of chromic anhydride (chromium trioxide, CrO3):[8]

K2Cr2O7 + 2 H2SO4 → 2 CrO3 + 2 KHSO4 + H2O

On heating with concentrated acid, oxygen is evolved:

2 K2Cr2O7 + 8 H2SO4 → 2 K2SO4 + 2 Cr2(SO4)3 + 8 H2O + 3 O2

Potassium dichromate is readily reduced by sulfur dioxide:[10]

K2Cr2O7 + H2SO4 + 3 SO2 → K2SO4 + Cr2(SO4)3 + H2O

In addition to providing a route to chromium(III) sulfate, this reaction was once the basis of a test for sulfur dioxide.

Potassium trichromate (K2Cr3O10) is prepared by evaporating a solution of potassium dichromate in nitric acid (d = 1.19) over concentrated sulfuric acid at low temperature. The tetrachromate (K2Cr4O13) is prepared by evaporating a solution of the trichromate in nitric acid (d = 1.4-1.5) slowly on a sand bath.[11]

Potassium dichromate reacts with hydrogen chloride to give chromyl chloride:[12]

K2Cr2O7 + 6 HCl → 2 CrO2Cl2 + 2 KCl + 3 H2O

Analytical reagent

Because it is non-hygroscopic, potassium dichromate was a common reagent in classical "wet tests" in analytical chemistry.

Organic chemistry

Potassium dichromate is an oxidising agent in organic chemistry. It is a member of a large collection of chromium-based oxidants.[13] Oxidations are often conducted in the presence of sulfuric acid, which favors the formation of chromium trioxide, which is the active oxidation reagent.

It is milder and more selective than potassium permanganate. It is often used interchangeably with sodium dichromate, although the use of chromates has declined owing to environmental concerns (see Hexavalent chromium). Especially in the presence of acids such as sulfuric acid, it is well known to oxidize alcohols. It converts primary alcohols into aldehydes[14] and, under more forcing conditions, into carboxylic acids. In contrast, potassium permanganate tends to give carboxylic acids as the sole products. Secondary alcohols are converted into ketones.[15] Tertiary alcohols are not oxidized.

Oxidations by dichromate are accompanied by a color change from yellow-orange for Cr(VI) to green for Cr(III). The color change, a colorimetric test, exhibited can be used as a test to distinguish aldehydes from ketones. Aldehydes are oxidized further, whereas ketones resist oxidation by dichromate salts.[16]

In some cases, potassium dichromate can cleave C=C bonds.[17]

Potassium dichromate is a classical reagent for the preparation of naphthoquinones.[18]

Leather

Potassium dichromate has few major applications, as the sodium salt is dominant industrially. The main use is as a precursor to potassium chrome alum, used in leather tanning.[8][19]

Wood treatment

Potassium dichromate is used to stain certain types of wood by darkening the tannins in the wood. It produces deep, rich browns that cannot be achieved with modern color dyes. It is a particularly effective treatment on mahogany.[20]

Natural occurrence

A ~10 mm crystal of potassium dichromate in the same form as the mineral lópezite

Potassium dichromate occurs naturally as the rare mineral lópezite. It has only been reported as vug fillings in the nitrate deposits of the Atacama Desert of Chile and in the Bushveld igneous complex of South Africa.[21]

Safety

Potassium dichromate is a prevalent allergen in patch tests (4.8%). Its presence in cement can cause contact dermatitis in construction workers after extended exposure.[22][23] In general, it is one of the most common causes of chromium dermatitis.[24] Aquatic organisms are vulnerable to poisoning by dichromate salts, but far less so than organic pollutants.[25]

As with other Cr(VI) compounds, potassium dichromate is carcinogenic.[26] The compound is also corrosive and exposure may damage eyes. Human exposure further causes impaired fertility.[5]

References

  1. "Potassium Dichromate Listing" (PDF). US EPA. 2015-07-23. Archived from the original (PDF) on June 28, 2011.
  2. Lide, David R., ed. (2004). CRC Handbook of Chemistry and Physics (85th ed.). Boca Raton, Florida: CRC Press. p. 4-76. ISBN 978-0-8493-0485-9.
  3. Binnewies, M.; Milke, E. (2002). Thermochemical Data of Elements and Compounds (2 ed.). Weinheim: Wiley-VCH. p. 405. ISBN 978-3-527-30524-7.
  4. Sigma-Aldrich Co., Potassium dichromate. Retrieved on 13 January 2026.
  5. "SDS - Potassium Dichromate". fishersci.com. ThermoFisher Scientific. 18 December 2025. Retrieved 13 January 2026.
  6. "PubChem: Potassium dichromate".
  7. "NIOSH Pocket Guide to Chemical Hazards".
  8. Anger, Gerd; Halstenberg, Jost; Hochgeschwender, Klaus; Scherhag, Christoph; Korallus, Ulrich; Knopf, Herbert; Schmidt, Peter; Ohlinger, Manfred. "Chromium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a07_067. ISBN 978-3-527-30673-2.
  9. Silber, C.; Breidenstein, B.; Heide, G.; Follner, H. (September 1999). "Growth Mechanism and Crystal Form of Potassium Dichromate". Crystal Research and Technology. 34 (8): 969–974. Bibcode:1999CryRT..34..969S. doi:10.1002/(SICI)1521-4079(199909)34:8<969::AID-CRAT969>3.0.CO;2-T.
  10. Hein, F.; Herzog, S. (2 December 2012) [1963]. "Chromium, Molybdenum, Tungsten, Uranium". In G. Brauer (ed.). Handbook of Preparative Inorganic Chemistry, 2nd Ed. Vol. 2. NY, NY: Academic Press. p. 1366. ISBN 9780323161299.
  11. Fedoroff, Basil T.; Sheffield, Oliver E. (1 January 1966). "C - Trichromates - Potassium trichromate & Tetrachromates - Potassium tetrachromate". Encyclopedia of Explosives and Related Items (PDF) (Technical report). Vol. 3, Chlorides through Detonating Relays. Picatinny Arsenal, NJ: U.S. Army Armament Research Development And Engineering Center - TACOM, ARDEC - Warheads, Energetics And Combat Support Armaments Center. pp. C288-9. AD0653029, PATR 2700.
  12. Sisler, Harry H. (2006) [1st pub. 1946]. Fernelius, W. Conard (ed.). Chromyl Chloride [Chromium(VI) Dioxychloride]. Inorganic Syntheses. Vol. 2 (1 ed.). Wiley. p. 205. doi:10.1002/9780470132333.ch63. ISBN 978-0-470-13161-9. Retrieved 2026-05-15.
  13. Cainelli, G.; Cardillo, G. (1981). Chromium Oxidations in Organic Chemistry. La Salle, IL: Open Court Publishers. p. 118–216.{{cite book}}: CS1 maint: multiple names: authors list (link)
  14. Charles D. Hurd and R. N. Meinert (1932). "Propionaldehyde". Organic Syntheses. 12: 64. doi:10.15227/orgsyn.012.0064.
  15. Luciano Lombardo (1987). "Methylenation of Carbonyl Compounds: (+)-3-Methylene-cis-p-Menthane". Organic Syntheses. 65: 81. doi:10.15227/orgsyn.065.0081.
  16. "oxidation of aldehydes and ketones". www.chemguide.co.uk. Retrieved 2025-10-30.
  17. Oliver Grummitt; Richard Egan; Allen Buck (1949). "Homophthalic Acid and Anhydride". Organic Syntheses. 29: 49. doi:10.15227/orgsyn.029.0049.
  18. Louis F. Fieser (1925). "1,4-Naphthoquinone". Organic Syntheses. 5: 79. doi:10.15227/orgsyn.005.0079.
  19. M. Saha; C. R. Srinivas; S. D. Shenoy; C. Balachandran (May 1993). "Footwear dermatitis". Contact Dermatitis. 28 (5): 260–264. doi:10.1111/j.1600-0536.1993.tb03428.x. PMID 8365123. S2CID 23159708.
  20. Jewitt, Jeff (1997). Hand-Applied Finishes. Newtown, Connecticut: Taunton Press. ISBN 978-1-56158-154-2.
  21. "Lópezite". mindat.org. Retrieved 13 January 2026.
  22. Sprung, Siegbert (2008). "Cement". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a05_489.pub2. ISBN 978-3-527-30673-2.
  23. Hedberg, Yolanda S.; Gumulka, Martin; Lind, Marie-Louise; Matura, Mihály; Lidén, Carola (2014). "Severe occupational chromium allergy despite cement legislation". Contact Dermatitis. 70 (5): 321–323. doi:10.1111/cod.12203. PMID 24731090.
  24. Bregnbak, David; Johansen, Jeanne D.; Jellesen, Morten S.; Zachariae, Claus; Menné, Torkil; Thyssen, Jacob P. (2015). "Chromium allergy and dermatitis: Prevalence and main findings". Contact Dermatitis. 73 (5): 261–280. doi:10.1111/cod.12436. PMID 26104877.
  25. Weltje, Lennart; Simpson, Peter; Gross, Melanie; Crane, Mark; Wheeler, James R. (2013). "Comparative Acute and Chronic Sensitivity of Fish and Amphibians: A Critical Review of Data". Environmental Toxicology and Chemistry. 32 (5): 984–994. Bibcode:2013EnvTC..32..984W. doi:10.1002/etc.2149. PMID 23381988.
  26. IARC (2012) [17-24 March 2009]. Volume 100C: Arsenic, Metals, Fibres, and Dusts (PDF). Lyon: International Agency for Research on Cancer. ISBN 978-92-832-0135-9. Archived from the original (PDF) on 2020-03-17. Retrieved 2020-01-05. There is sufficient evidence in humans for the carcinogenicity of chromium (VI) compounds. Chromium (VI) compounds cause cancer of the lung. Also positive associations have been observed between exposure to Chromium (VI) compounds and cancer of the nose and nasal sinuses. There is sufficient evidence in experimental animals for the carcinogenicity of chromium (VI) compounds. Chromium (VI) compounds are carcinogenic to humans (Group 1).