|Molar mass||134.45 g/mol (anhydrous)
170.48 g/mol (dihydrate)
|Appearance||yellow-brown solid (anhydrous)
blue-green solid (dihydrate)
|Density||3.386 g/cm3 (anhydrous)
2.51 g/cm3 (dihydrate)
|Melting point||498 °C (928 °F; 771 K) (anhydrous)
100 °C (dehydration of dihydrate)
|Boiling point||993 °C (1,819 °F; 1,266 K) (anhydrous, decomposes)|
|70.6 g/100 mL (0 °C)
75.7 g/100 mL (25 °C)
107.9 g/100 mL (100 °C)
68 g/100 mL (15 °C)
|Crystal structure||distorted CdI2 structure|
|EU classification||Not listed|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is: / ?)(|
Copper(II) chloride is the chemical compound with the chemical formula CuCl2. This is a light brown solid, which slowly absorbs moisture to form a blue-green dihydrate. The copper(II) chlorides are some of the most common copper(II) compounds, after copper sulfate.
- 1 Structure
- 2 Properties and reactions
- 3 Preparation
- 4 Uses
- 5 Safety
- 6 References
- 7 Further reading
- 8 External links
Anhydrous CuCl2 adopts a distorted cadmium iodide structure. In this motif, the copper centers are octahedral. Most copper(II) compounds exhibit distortions from idealized octahedral geometry due to the Jahn-Teller effect, which in this case describes the localization of one d-electron into a molecular orbital that is strongly antibonding with respect to a pair of chloride ligands. In CuCl2·2H2O, the copper again adopts a highly distorted octahedral geometry, the Cu(II) centers being surrounded by two water ligands and four chloride ligands, which bridge asymmetrically to other Cu centers.
Properties and reactions
Aqueous solution prepared from copper(II) chloride contain a range of copper(II) complexes depending on concentration, temperature, and the presence of additional chloride ions. These species include blue color of [Cu(H2O)6]2+ and yellow or red color of the halide complexes of the formula [CuCl2+x]x−.
Copper(II) hydroxide precipitates upon treating copper(II) chloride solutions with base:
- CuCl2 + 2 NaOH → Cu(OH)2 + 2 NaCl
Partial hydrolysis gives copper oxychloride, Cu2Cl(OH)3, a popular fungicide.
- 2 CuCl2 → 2 CuCl + Cl2
CuCl2 reacts with several metals to produce copper metal or copper(I) chloride with oxidation of the other metal. To convert copper(II) chloride to copper(I) derivatives, it can be convenient to reduce an aqueous solution with sulfur dioxide as the reductant:
- 2 CuCl2 + SO2 + 2 H2O → 2 CuCl + 2 HCl + H2SO4
2 + Cl−
2 + 2 Cl−
Some of these complexes can be crystallized from aqueous solution, and they adopt a wide variety of structures.
- CuCl2 + 2 C5H5N → [CuCl2(C5H5N)2] (tetragonal)
- CuCl2 + 2 (C6H5)3P=O → [CuCl2((C6H5)3P=O)2] (tetrahedral)
Copper(II) chloride is prepared commercially by the action of chlorination of copper:
- Cu + Cl2 + 2 H2O → CuCl2(H2O)2
Once prepared, a solution of CuCl2 may be purified by crystallization. A standard method takes the solution mixed in hot dilute hydrochloric acid, and causes the crystals to form by cooling in a CaCl2-ice bath.
There are indirect and rarely used means of using copper ions in solution to form copper(II) chloride. Electrolysis of aqueous sodium chloride with copper electrodes produces (among other things) a blue-green foam that can be collected and converted to the hydrate. While this is not usually done due to the emission of toxic chlorine gas, and the prevalence of the more general chloralkali process, the electrolysis will convert the copper metal to copper ions in solution forming the compound. Indeed any solution of copper ions can be mixed with hydrochloric acid and made into copper chloride by removing any other ions.
Copper(II) chloride occurs naturally as the very rare mineral tolbachite and the dihydrate eriochalcite. Both are found near fumaroles. More common are mixed oxyhydroxide-chlorides like atacamite Cu2(OH)3Cl, arising among Cu ore beds oxidation zones in arid climate (also known from some altered slags).
Co-catalyst in Wacker process
A major industrial application for copper(II) chloride is as a co-catalyst with palladium(II) chloride in the Wacker process. In this process, ethene (ethylene) is converted to ethanal (acetaldehyde) using water and air. During the reaction, PdCl2 is reduced to Pd, and the CuCl2 serves to re-oxidize this back to PdCl2. Air can then oxidize the resultant CuCl back to CuCl2, completing the cycle.
- C2H4 + PdCl2 + H2O → CH3CHO + Pd + 2 HCl
- Pd + 2 CuCl2 → 2 CuCl + PdCl2
- 4 CuCl + 4 HCl + O2 → 4 CuCl2 + 2 H2O
The overall process is:
- 2 C2H4 + O2 → 2 CH3CHO
Catalyst in production of chlorine
Copper(II) chloride is used as a catalyst in a variety of processes that produce chlorine by oxychlorination. The Deacon process takes place at about 400 to 450 °C in the presence of a copper chloride:
- 4 HCl + O2 → 2 Cl2 + 2 H2O
Copper(II) chloride is used in the Copper–chlorine cycle in which it splits steam into a copper oxygen compound and hydrogen chloride, and is later recovered in the cycle from the electrolysis of copper(I) chloride.
Other organic synthetic applications
Copper(II) chloride has some highly specialized applications in the synthesis of organic compounds. It effects 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:
CuCl2, in the presence of oxygen, can also oxidize phenols. The major product can be directed to give either a quinone or a coupled product from oxidative dimerization. The latter process provides a high-yield route to 1,1-binaphthol:
Such compounds are intermediates in the synthesis of BINAP and its derivatives
In humidity indicator cards (HICs), cobalt-free brown to azure (copper(II) chloride base) HICs can be found on the market. In 1998, the European Community (EC) classified items containing cobalt(II) chloride of 0.01 to 1% w/w as T (Toxic), with the corresponding R phrase of R49 (may cause cancer if inhaled). As a consequence, new cobalt-free humidity indicator cards have been developed that contain copper.
Copper(II) chloride can be toxic. Only concentrations below 5 ppm are allowed in drinking water by the US Environmental Protection Agency.
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