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Copper(II) chloride

Copper(II) chloride
IUPAC name Copper(II) chloride
Copper dichloride
Other names Cupric chloride
CAS number [7447-39-4] (ahydrous)
10125-13-0] (dihydrate)
RTECS number GL7000000
Molecular formula CuCl2
Molar mass 134.45 g/mol (anhydrous)
170.48 g/mol (dihydrate)
Appearance blue-green solid (dihydrate)
yellow-brown solid (anhydrous)
Density 3.386 g/cm³, solid
Melting point

100 °C (dehydration
of dihydrate)

Boiling point

decomposes at 993°C

Solubility in water 70.6 g/100 mL (0 °C)
75.7 g/100 mL (25 °C)
Solubility in ethanol 53 g/100 mL (15 °C)
Crystal structure distorted CdI2 structure
EU classification not listed
Flash point nonflammable
Related Compounds
Other anions Copper(II) fluoride
Copper(II) bromide
Copper(I) iodide
Other cations Copper(I) chloride
Silver chloride
Gold(III) chloride
Except where noted otherwise, data are given for
materials in their standard state
(at 25 °C, 100 kPa)

Infobox disclaimer and references

Copper(II) chloride is the chemical compound with the formula CuCl2. This a yellow-brown solid which slowly absorbs moisture to form a blue-green dihydrate. It occurs naturally as the very rare mineral eriochalcite.



Anhydrous CuCl2 adopts a distorted cadmium iodide structure. Most copper(II) compounds exhibit distortions from idealized octahedral geometry due to the Jahn-Teller effect, which in this case describes the localisation of one d-electron into a molecular orbital that is strongly antibonding with respect to a pair of ligands. In CuCl2(H2O)2 the copper can be described as a highly distorted octahedral complex, the Cu(II) center being surrounded by two water ligands and four chloride ligands, which bridge asymmetrically to other Cu centers.[1]


Copper(II) chloride dissociates in aqueous solution to give the blue color of [Cu(H2O)6]2+ and yellow or red color of the halide complexes of the formula [CuCl2+x]x-. Concentrated solutions of CuCl2 appear green because of the combination of these various chromophores.

It is a weak Lewis acid, and a mild oxidising agent. It has a crystal structure consisting of polymeric chains of flat CuCl4 units with opposite edges shared. It decomposes to CuCl and Cl2 at 1000 °C:

2 CuCl2(s) → 2 CuCl(s) + Cl2(g)

In its reaction with HCl (or other chloride sources) to form the complex ions CuCl3- and CuCl42-.[2]

Some of these complexes can be crystallized from aqueous solution, and they adopt a wide variety structural types (Fig. 1).

Copper(II) chloride also forms a rich variety of other coordination complexes with ligands such as pyridine or triphenylphosphine oxide:

CuCl2 + 2 C5H5N → [CuCl2(C5H5N)2] (tetragonal)
CuCl2 + 2 (C6H5)3P=O → [CuCl2((C6H5)3P=O)2] (tetrahedral)

However "soft" ligands such as phosphines (e.g., triphenylphosphine), iodide, and cyanide as well as some tertiary amines cause reduction to give copper(I) complexes. To convert copper(II) chloride to copper(I) derivatives it is generally more convenient to reduce an aqueous solution with the reducing agent sulfur dioxide:

2 CuCl2(aq) + SO2 → 2 CuCl(s) + 2 HCl(aq) + H2SO4(aq)

CuCl2 can simply react as a source of Cu2+ in precipitation reactions for making insoluble copper(II) salts, for example copper(II) hydroxide, which can then decompose above 30 °C to give copper(II) oxide:

CuCl2(aq) + 2 NaOH(aq) → Cu(OH)2(s) + 2 NaCl(aq)

Followed by

Cu(OH)2(s) → CuO(s) + H2O(l)


Copper(II) chloride is prepared by the action of hydrochloric acid on copper(II) oxide, copper(II) hydroxide or copper(II) carbonate, for example:

CuO(s) + 2 HCl(aq) → CuCl2(aq) + H2O(l) Anhydrous CuCl2 may be prepared directly by union of the elements, copper and chlorine.

CuCl2 may be purified by crystallisation from hot dilute hydrochloric acid, by cooling in a CaCl2-ice bath[7].


A major industrial application for copper(II) chloride is as a co-catalyst (along with palladium(II) chloride) in the Wacker process. In this process, ethene (ethylene) is converted to ethanal (acetaldehyde) using water and air. In the process PdCl2 is reduced to Pd, and the CuCl2 serves to re-oxidise this back to PdCl2. Air can then oxidise the resultant CuCl back to CuCl2, completing the cycle.

(1) C2H4(g) + PdCl2(aq) + H2O (l) → CH3CHO (aq) + Pd(s) + 2 HCl(aq)

(2) Pd(s) + 2 CuCl2(aq) → 2 CuCl(s) + PdCl2(aq)

(3) 2 CuCl(s) + 2 HCl(aq) +1/2O2(g) → 2 CuCl2(aq) + H2O(l)

Overall process: C2H4 +1/2O2 → CH3CHO

Copper(II) chloride has a variety of applications in organic synthesis[7]. It can effect chlorination of aromatic hydrocarbons- this is often performed in the presence of aluminium oxide. It is able to chlorinate the alpha position of carbonyl compounds[8]:

This reaction is performed in a polar solvent such as DMF, often in the presence of lithium chloride, which speeds up the reaction rate.

CuCl2, in the presence of oxygen, can also oxidise phenols. The major product can be directed to give either a quinone or a coupled product from oxidative dimerisation. The latter process provides a high-yield synthesis of 1,1-binaphthol (also called BINOL) and its derivatives, these can even be made as a single enantiomer in high enantiomeric excess[9]:

Such compounds are valuable intermediates in the synthesis of BINAP and its derivatives, popular as chiral ligands for asymmetric hydrogenation catalysts.

CuCl2 also catalyses the free radical addition of sulfonyl chlorides to alkenes; the alpha-chlorosulfone may then undergo elimination with base to give a vinyl sulfone product.

Copper(II) chloride is also used in pyrotechnics as a blue/green coloring agent.


Although copper is an essential element, all metal salts are potentially toxic if mishandled. See MSDS.


  1. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  2. ^ Gill, N. S.; Taylor, F. B., "Tetrahalo Complexes of Dipositive Metals in the First Transition Series", Inorganic Syntheses, 1967, volume 9, pages 136-142.
  1. Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. 
  2. Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  3. The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
  4. D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
  5. A. F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
  6. J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
  7. S. H. Bertz, E. H. Fairchild, in Handbook of Reagents for Organic Synthesis, Volume 1: Reagents, Auxiliaries and Catalysts for C-C Bond Formation, (R. M. Coates, S. E. Denmark, eds.), pp. 220-3, Wiley, New York, 1999.
  8. C. E. Castro, E. J. Gaughan, D. C. Owsley, Journal of Organic Chemistry, 30, 587 (1965).
  9. J. Brussee, J. L. G. Groenendijk, J. M. Koppele, A. C. A. Jansen, Tetrahedron, 41, 3313 (1985).
  10. Fieser & Fieser Reagents for Organic Synthesis Volume 5, p158, Wiley, New York, 1975.
  11. D. W. Smith (1976). "Chlorocuprates(II)". Coordination Chemistry Reviews 21 (2-3): 93-158. doi:10.1016/S0010-8545(00)80445-2.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Copper(II)_chloride". A list of authors is available in Wikipedia.
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