Copper(II) chloride

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Copper(II) chloride

General
Systematic name	Copper(II) chloride Copper dichloride
Other names	Cupric chloride
Molecular formula	CuCl₂
Molar mass	134.45 g/mol (anhydrous)
Appearance	blue-green crystalline solid (dihydrate) yellow-brown powder (anhydrous)
CAS number	[7447-39-4]
Properties
Density and phase	3.386 g/cm³, solid
Solubility in water	70.6 g/100 ml (0 °C) 75.7 g/100 ml (25 °C)
Solubility in methanol	68 g/100 ml (15 °C)
Solubility in ethanol	53 g/100 ml (15 °C)
Melting point	620 °C (anhydrous)
Boiling point	993 °C
Structure
Coordination geometry	Octahedral
Crystal structure	CdI₂ structure
Dipole moment	? D
Hazards
MSDS	External MSDS
EU classification	not listed
NFPA 704
Flash point	non flammable
RTECS number	?
Supplementary data page
Structure and properties	n, ε_r, etc.
Thermodynamic data	Phase behaviour Solid, liquid, gas
Spectral data	UV, IR, NMR, MS
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 higher chloride of copper, with the formula CuCl₂. It occurs naturally as the mineral eriochalcite. It is a yellow-brown solid which slowly absorbs moisture to form a blue-green dihydrate.

It is ionic and highly soluble in water. Chemically it behaves as a weak Lewis acid, and under certain conditions it can act as 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.

1 Chemical Properties
2 Preparation
3 Uses
4 Precautions
5 References
6 External links

[edit] Chemical Properties

Copper(II) chloride is an ionic compound which dissociates in aqueous solution, but coordination of chloride to Cu²⁺ does partially occur. This means that concentrated solutions of CuCl₂ are green- a combination of the blue color of [Cu(H₂O)₆]²⁺ with the yellow or red color of the halide complexes.

CuCl₂ also behaves as a mild Lewis acid, for example in its reaction with HCl (or other chloride sources) to form the complex ions CuCl₃^- and CuCl₄^2-.

Some of these complexes can be isolated as crystals from aqueous solution, and they form 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 certain other ligands such as phosphines (e.g., triphenylphosphine) and even some tertiary amines cause reduction to copper(I) complexes.

Reduction to copper(I) chloride can be effected simply by heating CuCl₂ at high temperatures (about 1000 °C):

2 CuCl₂(s) → 2 CuCl(s) + Cl₂(g)

However, it is generally more convenient to work in aqueous solution, and to use a reducing agent such as sulfur dioxide to make CuCl:

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)

Then Cu(OH)₂(s) → CuO(s) + 2 H₂O(l)

[edit] Preparation

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].

CuCl₂ is also produced when a penny is put in a quantity of household chlorine bleach.

[edit] Uses

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(g) + ¹/₂O₂(g) → CH₃CHO (aq)

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 green colouring agent.

[edit] Precautions

Copper salts are toxic, and can be fatal. Wear gloves and goggles, avoid ingestion or inhalation.

[edit] References

Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.). Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
D. Nicholls, Complexes and First-Row Transition Elements, Macmillan Press, London, 1973.
A. F. Wells, 'Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984.
J. March, Advanced Organic Chemistry, 4th ed., p. 723, Wiley, New York, 1992.
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.
C. E. Castro, E. J. Gaughan, D. C. Owsley, Journal of Organic Chemistry, 30, 587 (1965).
J. Brussee, J. L. G. Groenendijk, J. M. Koppele, A. C. A. Jansen, Tetrahedron, 41, 3313 (1985).
Fieser & Fieser Reagents for Organic Synthesis Volume 5, p158, Wiley, New York, 1975.