Copper(II) chlorate

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Copper(II) chlorate
IUPAC name
Copper(2+) chlorate hydrate (1:2:4)
Other names
Copper(II) chlorate; Cupric chlorate
3D model (JSmol)
ECHA InfoCard 100.035.228 Edit this at Wikidata
EC Number
  • 238-767-0
UN number 2721
  • InChI=1S/2ClHO3.Cu.4H2O/c2*2-1(3)4;;;;;/h2*(H,2,3,4);;4*1H2/q;;+2;;;;/p-2
  • InChI=1S/2ClHO3.Cu/c2*2-1(3)4;/h2*(H,2,3,4);/q;;+2/p-2
  • O.O.O.O.[O-]Cl(=O)=O.[O-]Cl(=O)=O.[Cu+2]
Molar mass 302.509
Appearance Light blue
Density 2.26 g cm−3
Melting point 73 °C (anhydrous)
65 °C (hexahydrate)
Boiling point decomposes
highly water-soluble
54.59 g/100 mL (-31 °C)
57.12 g/100 mL (-21 °C)
58.51 g/100 mL (0.8 °C)
62.17 g/100 mL (18 °C)
66.17 g/100 mL (45 °C)
69.42 g/100 mL (59.6 °C)
76.9 g/100 mL (71 °C)
141 g/100 mL (0 °C)
164.4 g/100 mL (18 °C)
195.6 g/100 mL (45 °C)
332 g/100 mL (70 °C)
Solubility soluble in acetone and ethanol (hexahydrate)
a = 12.924 Å, b = 9.502 Å, c = 7.233 Å
880.4 Å3
distorted octahedral
Occupational safety and health (OHS/OSH):
Main hazards
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Copper(II) chlorate is a chemical compound of the transition metal copper and the chlorate anion with basic formula Cu(ClO3)2. Copper chlorate is an oxidiser.[5] It commonly forms the tetrahydrate, Cu(ClO3)2·4H2O.


Copper chlorate can be made by combining a hot one molar solution of copper sulfate, with barium chlorate, which results in the precipitation of barium sulfate. When the solution is filtered, cooled and evaporated under a vacuum blue crystals form.[6]

CuSO4 + Ba(ClO3)2 Cu(ClO3)2 + BaSO4(s)


In 1902, A. Meusser investigated solubility of copper chlorate and found that it melted and started decomposing above 73 °C, giving off chlorine.[7]

Copper chlorate decomposes when heated, giving off a yellow gas, which contains chlorine, oxygen and chlorine dioxide.[8] A green solid is left that is a basic copper salt.[9]

2 Cu(ClO3)2 2 CuO + Cl2 + 3 O2 + 2 ClO2

Sulfur is highly reactive with copper chlorate, and it is important not to cross contaminate these chemicals, for example in pyrotechnic making.[10]


Copper(II) chlorate commonly crystallizes as a tetrahydrate, though a hexahydrate is also known. Tetraaquacopper(II) chlorate, Cu(ClO3)2·4H2O, has an orthorhombic crystal structure.[4] Each copper atom is octahedrally coordinated, surrounded by four oxygen atoms of water, and two oxygen atoms from chlorate groups, which are opposite each other. Water is closer to the copper than chlorate, 1.944 Å compared to 2.396 Å, exhibiting the Jahn-Teller effect. The chlorate groups take the shape of a distorted tetrahedron. At 298 K (25 °C), the chlorine-oxygen distances in each chlorate ion are 1.498, 1.488 and 1.468 Å, with the longest being the oxygen next to copper. The ∠O-Cu-O (angle subtended at copper by oxygen atoms) is 105.2°, 108.3°, and 106.8°. At lower temperatures (233 K, −40 °C), the water molecules and copper-chlorate distance shrink.[4]


François-Marie Chertier used tetraamminecopper(II) chlorate to colour flames blue in 1843. This material was abbreviated TACC with formula Cu(NH3)4(ClO3)2. TACC explodes on impact.[11]

The substance became known as Chertier'c copper for use in blue coloured pyrotechnics.[12] However its deliquescence causes a problem.[13] Mixtures with other metal salts can yield violet or lilac colours also.[14]

It has also been used to colour copper brown.[15]


  1. ^ Seidell, Atherton (1919). A. Solubilities of inorganic and organic compounds. - 3ed., vol.1 (PDF) (2 ed.). New York: D. Van Nostrand Company. p. 264. Archived from the original (PDF) on 2018-02-03. Retrieved 2018-02-01.
  2. ^ Woolley, E. M.; Miyamoto, H.; Salomon, M. (1990). Copper and Silver Halates (PDF). Elsevier. ISBN 9781483286051.
  3. ^ "copper(II) chlorate hexahydrate".
  4. ^ a b c Blackburn, A. C.; Gallucci, J. C.; Gerkin, R. E. (1 August 1991). "Structure of tetraaquacopper(II) chlorate at 296 and 223 K". Acta Crystallographica Section B. B47 (4): 474–479. doi:10.1107/S0108768191000435. ISSN 0108-7681. PMID 1930830.
  5. ^ Lewis, Richard J. (2008). Hazardous Chemicals Desk Reference. John Wiley & Sons. p. 384. ISBN 9780470334454.
  6. ^ Suhara, Masahiko (April 1973). "The Temperature Dependence of the Nuclear Quadrupole Resonance of 35Cl in KClO3, AgClO3, Ba(ClO3)2·H2 O, and Cu(ClO3) 2·6H2O". Bulletin of the Chemical Society of Japan. 46 (4): 1053–1055. doi:10.1246/bcsj.46.1053.
  7. ^ Meusser, A. (April 1902). "Metallchlorate. Studien über die Löslichkeit der Salze. X". Berichte der Deutschen Chemischen Gesellschaft. 35 (2): 1414–1424. doi:10.1002/cber.19020350240.
  8. ^ Rosenstiehl, A. (September 1876). "The Theory of Formation of Aniline Black". Journal of the Chemical Society. London. 30 (165): 311.
  9. ^ Waechter, M. Alexander (30 April 2009). "On the preparation and properties of certain chlorates". Philosophical Magazine. 3rd Series. 25 (165): 235–237. doi:10.1080/14786444408644978.
  10. ^ Bretherick, L. (1990). Bretherick's Handbook of Reactive Chemical Hazards. Butterworths. p. 975. ISBN 9780750601030.
  11. ^ Kosanke, K. L.; Sturman, Barry T.; Winokur, Robert M.; Kosanke, B. J. (2012). Encyclopedic Dictionary of Pyrotechnics: (and Related Subjects). Journal of Pyrotechnics. p. 1107. ISBN 9781889526218.
  12. ^ Browne, W. H. (1873). The art of pyrotechny. London: The Bazaar. p. 35.
  13. ^ Thorpe, Sir Thomas Edward (1924). A Dictionary of Applied Chemistry. Longmans, Green, and Company. p. 520.
  14. ^ Hiscox, G. D. (1931). Henley's twentieth century formulas, recipes and processes. Рипол Классик. pp. 609–610. ISBN 9785876347008.
  15. ^ Krause, Hugo (1938). Metal coloring and finishing: latest practical methods for coloring and finishing metals of all kinds. Chemical publishing co. of N. Y., inc. p. 96.