Potassium chromate

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Potassium chromate
Potassium chromate sample.jpg
Potassium chromate sample on a watchglass.
Names
IUPAC name
Potassium chromate
Other names
Chromic acid, dipotassium salt
Dipotassium monochromate
Dipotassium chromate
Tarapacaite
Identifiers
Jmol-3D images Image
Properties
K2CrO4
Molar mass 194.19 g/mol
Appearance Yellow solid
Odor Odorless
Density 2.732 g/cm3
Melting point 968 °C (1,774 °F; 1,241 K)
Boiling point 1,000 °C (1,830 °F; 1,270 K)
56.3 g/100 ml (0 °C)
60.0 g/100 ml (10 °C)
63.7 g/100 ml (20 °C)
67.8 g/100 ml (40 °C)
70.1 g/100 ml (60 °C)
79.2 g/100 ml (100 °C)
Solubility Soluble in ethylene glycol, trifluoroacetic acid
Slightly soluble in acetic anhydride, thionyl chloride
Insoluble in alcohols, alkanes, liq. sulfur dioxide
Vapor pressure ~0 mmHg
Thermochemistry
-1,383.27 kJ/kmol
Hazards
Safety data sheet ScienceLab
Related compounds
Related compounds
Potassium dichromate
Potassium ferrate
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
Infobox references

Potassium chromate (K2CrO4) is the potassium salt of the chromate anion. It is a yellow solid, soluble in water. As a Cr(VI) compound, it is a strong carcinogen and an oxidizer.

Properties

Chemical

The addition of an acid to an aqueous solution of potassium chromate will yield potassium dichromate.

2 K2CrO4 + H2SO4 → K2Cr2O7 + K2SO4 + H2O

Other acids, such as boric acid, hydrochloric acid, nitric acid, glacial acetic acid, oxalic acid, phosphoric acid can also be used. Carbon dioxide also works.

Physical

Potassium chromate is an odorless yellow solid, soluble in water, trifluoroacetic acid, liq. hydrogen cyanide, phosphoryl chloride. It is insoluble in acetone, anhydrous ammonia, aniline, benzonitrile, liq. chlorine, ethyl acetate, methyl acetate, molten antimony(III) chloride, vanadyl chloride (both cold and hot).

It has an average density of 2.7320 g/cm3.

Availability

Potassium chromate is sold by various chemical suppliers. It can also be found on eBay.

It occurs in nature as the rare mineral tarapacaite, which can be found in some nitrate ores from the Tarapacá province, Chile, in the arid Atacama desert.

In some countries, the sale of Cr(VI) compounds is regulated.

Preparation

Potassium chromate can be made by adding potassium hydroxide and conc. hydrogen peroxide or bleach to a suspension of chromium(III) hydroxide.

Cr(OH)3 + 2 KOH + H2O2 → K2CrO4 + H2O + 3/2 O2[1]
2 Cr(OH)3 + 6 KClO → 2 K2CrO4 + 4 KCl + 2 Cl2

The chromium hydroxide can be made by dissolving stainless steel in hydrochloric acid, followed by precipitation with sodium carbonate.[2]

If you're using common household bleach, you will end up with sodium chromate. To convert it into potassium chromate, you will need to add a potassium salt, such as potassium carbonate, filter the resulting solution and then recrystallize it.

Chromium(IV) oxide can also replace chromium(III) hydroxide.[3]

It can also be made by adding potassium hydroxide to potassium dichromate in solution.

Heating potassium dichromate will result in potassium chromate, chromium(III) oxide and oxygen.

4 K2Cr2O7 → 4 K2CrO4 + 2 Cr2O3 + 3 O2

Heating a mixture of potassium nitrate and chromium(III) oxide with traces of water will also give potassium chromate.[4] This reaction however does not work with indurated chromium oxide (like the OTC chromium oxide pigment).

Projects

Handling

Safety

Potassium chromate, like all hexavalent chromium compounds, is highly toxic and carcinogenic on ingestion or inhalation. Handling it without gloves can cause dermatitis, and can also be absorbed through the skin in small amounts, usually if wet. Aqueous solutions are notorious for staining most materials.

Always wear gloves and goggles when handling it, and a dust mask or respirator when handling it as a powder to avoid inhalation of it, which could be fatal.

Storage

Potassium chromate should be stored in closed bottles, with a visible label and a hazard symbol.

Disposal

Potassium chromate can be safely reduced to the less harmful Cr(III) oxide with a reducing agent, such as ascorbic acid, or various sulfites/metabisulfites/tiosulfates.

References

  1. Martinon; Bulletin de la Societe Chimique de France; vol. 45; (1886); p. 862 - 864
  2. https://www.youtube.com/watch?v=F_W-IyUTM5M
  3. Martinon; Bulletin de la Societe Chimique de France; vol. 45; (1886); p. 862 - 864
  4. Guignet; Bulletin de la Societe Chimique de France; (1859); p. 9 - 11

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