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Copper(II) chloride

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Copper(II) chloride

Anhydrous

Anhydrous

Dihydrate
Names
IUPAC names
Copper(II) chloride
Copper dichloride
Other names
Cupric chloride
Identifiers
3D model (JSmol)
ChEBI
ChEMBL
ChemSpider
ECHA InfoCard 100.028.373 Edit this at Wikidata
RTECS number
  • GL7000000
UNII
  • InChI=1S/2ClH.Cu/h2*1H;/q;;+2/p-2 checkY
    Key: ORTQZVOHEJQUHG-UHFFFAOYSA-L checkY
  • InChI=1/2ClH.Cu/h2*1H;/q;;+2/p-2/rCl2Cu/c1-3-2
    Key: ORTQZVOHEJQUHG-LRIOHBSEAE
  • InChI=1/2ClH.Cu/h2*1H;/q;;+2/p-2
    Key: ORTQZVOHEJQUHG-NUQVWONBAE
  • Cl[Cu]Cl
  • [Cu+2].[Cl-].[Cl-]
Properties
CuCl2
Molar mass 134.45 g/mol (anhydrous)
170.48 g/mol (dihydrate)
Appearance yellow-brown solid (anhydrous)
blue-green solid (dihydrate)
Odor odorless
Density 3.386 g/cm3 (anhydrous)
2.51 g/cm3 (dihydrate)
Melting point 498 °C (anhydrous)
100 °C (dehydration of dihydrate)
Boiling point 993 °C (anhydrous, decomp)
70.6 g/100 mL (0 °C)
75.7 g/100 mL (25 °C)
107.9 g/100 mL (100 °C)
Solubility methanol:
68 g/ 100 mL (15 °C)
ethanol:
53 g/100 mL (15 °C)
soluble in acetone
Structure
distorted CdI2 structure
Octahedral
Hazards
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 0: Will not burn. E.g. waterInstability 1: Normally stable, but can become unstable at elevated temperatures and pressures. E.g. calciumSpecial hazards (white): no code
2
0
1
Flash point Non-flammable
Related compounds
Other anions
 Copper(II) fluoride
Copper(II) bromide
Other cations
 Copper(I) chloride
Silver chloride
Gold(III) chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
☒N verify (what is checkY☒N ?)

Copper(II) chloride is the chemical compound with the 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.

Structure

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.[1]

Copper(II) chloride is paramagnetic. Of historical interest, CuCl2·2H2O was used in the first electron paramagnetic resonance measurements by Yevgeny Zavoisky in 1944.[2][3]

Properties and reactions

Aqueous solutions of "copper(II) chloride". Greenish when high on [Cl–], more blue when lower on [Cl–].

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−.[4]

Copper(II) chloride dihydrate crystal

It decomposes to CuCl and Cl2 at 1000 °C:

2 CuCl2 → 2 CuCl + Cl2

It reacts with HCl or other chloride sources to form complex ions: the red CuCl3−, and the yellow CuCl42−.[5]

CuCl
2 + 2 Cl−
 ⇌ CuCl−
3 + Cl−
 ⇌ CuCl2−
4

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

Copper(II) hydroxide precipitates upon treating copper(II) chloride solutions with base:

CuCl2 + 2 NaOH → Cu(OH)2 + 2 NaCl

Copper(II) chloride also forms a variety of coordination complexes with ligands such as pyridine and 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 sulfur dioxide as the reductant:

2 CuCl2 + SO2 + 2 H2O → 2 CuCl + 2 HCl + H2SO4

Hydrolysis give the copper oxychloride, Cu2Cl(OH)3, a popular fungicide.

Preparation

Copper(II) chloride is prepared commercially by the action of chlorination of copper:

Cu + Cl2 + 2 H2O → CuCl2(H2O)2

It can also be generated by treatment of the hydroxide, oxide, or copper(II) carbonate with hydrochloric acid. 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.

Anhydrous CuCl2 may be prepared directly by union of the elements, copper and chlorine.

CuCl2 may be purified by crystallization from hot dilute hydrochloric acid, by cooling in a CaCl2-ice bath.[6][7]

Natural occurrence

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).

Uses

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.

  1. C2H4 + PdCl2 + H2O → CH3CHO + Pd + 2 HCl
  2. Pd + 2 CuCl2 → 2 CuCl + PdCl2
  3. 4 CuCl + 4 HCl + O2 → 4 CuCl2 + 2 H2O

The overall process is:

2 C2H4 + O2 → 2 CH3CHO

Chlorinations

Copper(II) chloride catalyzes the chlorination in the production of vinyl chloride and dichloroethane.[8]

Other organic synthetic applications

Copper(II) chloride has a variety of specialized applications in the synthesis of organic compounds.[6] 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:[9]

Alpha chlorination of an aldehyde using CuCl2.

This reaction is performed in a polar solvent such as dimethylformamide (DMF), often in the presence of lithium chloride, which accelerates the reaction.

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:[10]

Coupling of beta-naphthol using CuCl2.

Such compounds are intermediates in the synthesis of BINAP and its derivatives

Copper(II) chloride dihydrate promotes the hydrolysis of acetonides, i.e., for deprotection to regenerate diols[11] or aminoalcohols, as in this example (where TBDPS = tert-butyldiphenylsilyl):[12]

Deprotection of an acetonide using CuCl2·2H2O.

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.[citation needed]

Niche uses

Copper(II) chloride is also used in pyrotechnics as a blue/green coloring agent. In a flame test, copper chlorides, like all copper compounds, emit green-blue.

Safety

It is toxic and only concentrations below 5 ppm are allowed in drinking water by the US Environmental Protection Agency.

References

  1. ^ Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN 0-19-855370-6.
  2. ^ Peter Baláž (2008). Mechanochemistry in Nanoscience and Minerals Engineering. Springer. p. 167. ISBN 3-540-74854-7.
  3. ^ Marina Brustolon (2009). Electron paramagnetic resonance: a practitioner's toolkit. John Wiley and Sons. p. 3. ISBN 0-470-25882-9.
  4. ^ Greenwood, N. N. and Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
  5. ^ Naida S. Gill; et al. (1967). "Inorganic Syntheses". Inorg. Synth. Inorganic Syntheses. 9: 136–142. doi:10.1002/9780470132401.ch37. ISBN 978-0-470-13240-1. {{cite journal}}: |chapter= ignored (help); Explicit use of et al. in: |author= (help)
  6. ^ a b 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.
  7. ^ W. L. F. Armarego, Christina Li Lin Chai (2009-05-22). Purification of Laboratory Chemicals (Google Books excerpt) (6th ed.). Butterworth-Heinemann. p. 461. ISBN 1-85617-567-7.
  8. ^ H.Wayne Richardson, "Copper Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim, doi:10.1002/14356007.a07_567
  9. ^ C. E. Castro, E. J. Gaughan, D. C. Owsley (1965). "Cupric Halide Halogenations". Journal of Organic Chemistry. 30 (2): 587. doi:10.1021/jo01013a069.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  10. ^ J. Brussee, J. L. G. Groenendijk, J. M. Koppele, A. C. A. Jansen (1985). "On the mechanism of the formation of s(−)-(1, 1'-binaphthalene)-2,2'-diol via copper(II)amine complexes". Tetrahedron. 41 (16): 3313. doi:10.1016/S0040-4020(01)96682-7.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  11. ^ Chandrasekhar, M. (2003). "Total Synthesis of (+)-Boronolide, (+)-Deacetylboronolide, and (+)-Dideacetylboronolide". Journal of Organic Chemistry. 68 (10): 4039–4045. doi:10.1021/jo0269058. PMID 12737588. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)
  12. ^ Krishna, Palakodety Radha (2007). "A stereoselective total synthesis of (-)-andrachcinidine via an olefin cross-metathesis protocol". Tetrahedron Letters. 48 (41). Elsevier: 7279–7282. doi:10.1016/j.tetlet.2007.08.053. {{cite journal}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)

Further reading

  1. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  2. Lide, David R. (1990). CRC handbook of chemistry and physics: a ready-reference book of chemical and physical data. Boca Raton: CRC Press. ISBN 0-8493-0471-7.
  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. Fieser & Fieser Reagents for Organic Synthesis Volume 5, p158, Wiley, New York, 1975.
  8. D. W. Smith (1976). "Chlorocuprates(II)". Coordination Chemistry Reviews. 21 (2–3): 93–158. doi:10.1016/S0010-8545(00)80445-2.

External links

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