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Copper(II) chloride

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Copper(II) chloride

Anhydrous
  Copper, Cu
  Chlorine, Cl

Anhydrous

Dihydrate
Names
IUPAC name
Copper(II) chloride
Other names
Cupric chloride
Identifiers
3D model (JSmol)
8128168
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.028.373 Edit this at Wikidata
EC Number
  • 231-210-2
9300
RTECS number
  • GL7000000
UNII
UN number 2802
  • 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 630 °C (1,166 °F; 903 K) (extrapolated)
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)
Solubility methanol:
68 g/(100 mL) (15 °C)


ethanol:
53 g/(100 mL) (15 °C)
soluble in acetone

+1080·10−6 cm3/mol
Structure[1][2]
monoclinic (β = 121°) (anhydrous)
orthorhombic (dihydrate)
C2/m (anhydrous)
Pbmn (dihydrate)
a = 6.85 Å (anhydrous)
7.41 Å (dihydrate), b = 3.30 Å (anhydrous)
8.09 Å (dihydrate), c = 6.70 Å (anhydrous)
3.75 Å (dihydrate)
Octahedral
Hazards
GHS labelling:
GHS05: CorrosiveGHS06: ToxicGHS07: Exclamation markGHS09: Environmental hazard
Danger
H301, H302, H312, H315, H318, H319, H335, H410, H411
P261, P264, P270, P271, P273, P280, P301+P310, P301+P312, P302+P352, P304+P340, P305+P351+P338, P310, P312, P321, P322, P330, P332+P313, P337+P313, P362, P363, P391, P403+P233, P405, P501
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
NIOSH (US health exposure limits):
PEL (Permissible)
 TWA 1 mg/m3 (as Cu)[3]
REL (Recommended)
 TWA 1 mg/m3 (as Cu)[3]
IDLH (Immediate danger)
 TWA 100 mg/m3 (as Cu)[3]
Safety data sheet (SDS) Fisher Scientific
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, also known as cupric chloride, is an inorganic compound with the chemical formula CuCl2. The monoclinic yellowish-brown anhydrous form slowly absorbs moisture to form the orthorhombic blue-green dihydrate CuCl2·2H2O, with two water molecules of hydration. It is industrially produced for use as a co-catalyst in the Wacker process.

Both the anhydrous and the dihydrate forms occur naturally as the rare minerals tolbachite and eriochalcite, respectively.

Structure

Structure of copper(II) chloride dihydrate
  Copper, Cu
  Oxygen, O
  Chlorine, Cl
  Hydrogen, H

Anhydrous copper(II) chloride adopts a distorted cadmium iodide structure. In this structure, 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.[4][5]

Copper(II) chloride is paramagnetic. Of historical interest, CuCl2·2H2O was used in the first electron paramagnetic resonance measurements by Yevgeny Zavoisky in 1944.[6][7]

Properties and reactions

Aqueous solutions of copper(II) chloride. Greenish when high in Cl−, more blue when lower in Cl−.

Aqueous solutions 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 the blue color of [Cu(H2O)6]2+ and the yellow or red color of the halide complexes of the formula [CuCl2+x]x−.[5]

Hydrolysis

When copper(II) chloride solutions are treated with a base, a precipitation of copper(II) hydroxide occurs:[8]

CuCl2 + 2 NaOH → Cu(OH)2 + 2 NaCl

Partial hydrolysis gives dicopper chloride trihydroxide, Cu2(OH)3Cl, a popular fungicide.[8] When an aqueous solution of copper(II) chloride is left in the air and isn't stabilized by a small amount of acid, it is prone to undergo slight hydrolysis.[5]

Redox and decomposition

Copper(II) chloride is a mild oxidant. It starts to decompose to copper(I) chloride and chlorine gas around 400 °C (752 °F) and is completely decomposed near 1,000 °C (1,830 °F):[8][9][10][11]

2 CuCl2 → 2 CuCl + Cl2

The reported melting point of copper(II) chloride of 498 °C (928 °F) is a melt of a mixture of copper(I) chloride and copper(II) chloride. The true melting point of 630 °C (1,166 °F) can be extrapolated by using the melting points of the mixtures of CuCl and CuCl2.[12][13] Copper(II) chloride reacts with several metals to produce copper metal or copper(I) chloride (CuCl) with oxidation of the other metal. To convert copper(II) chloride to copper(I) chloride, it can be convenient to reduce an aqueous solution with sulfur dioxide as the reductant:[8]

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

Coordination complexes

CuCl2 reacts with HCl or other chloride sources to form complex ions: the red [CuCl3]− (found in potassium trichloridocuprate(II) K[CuCl3]) (it is a dimer in reality, [Cu2Cl6]2−, a couple of tetrahedrons that share an edge), and the green or yellow [CuCl4]2− (found in potassium tetrachloridocuprate(II) K2[CuCl4]).[5][14][15]

CuCl2 + Cl− ⇌ [CuCl3]−
CuCl2 + 2 Cl− ⇌ [CuCl4]2−

Some of these complexes can be crystallized from aqueous solution, and they adopt a wide variety of structures.[14]

Copper(II) chloride also forms a variety of coordination complexes with ligands such as ammonia, pyridine and triphenylphosphine oxide:[8][5][16]

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 induce reduction to give copper(I) complexes.[5]

Preparation

Copper(II) chloride is prepared commercially by the action of chlorination of copper. Copper at red heat (300-400°C) combines directly with chlorine gas, giving (molten) copper(II) chloride. The reaction is very exothermic.[8][15]

Cu(s) + Cl2(g) → CuCl2(l)

A solution of copper(II) chloride is commercially produced by adding chlorine gas to a circulating mixture of hydrochloric acid and copper. From this solution, the dihydrate can be produced by evaporation.[8][10]

Although copper metal itself cannot be oxidized by hydrochloric acid, copper-containing bases such as the hydroxide, oxide, or copper(II) carbonate can react to form CuCl2 in an acid-base reaction which can subsequently be heated above 100 °C (212 °F) to produce the anhydrous derivative.[8][10]

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 calcium chloride (CaCl2) ice bath.[17][18]

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 a copper chloride by removing any other ions.[19]

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.[20]

  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:[20]

2 C2H4 + O2 → 2 CH3CHO

In organic synthesis

Copper(II) chloride has some highly specialized applications in the synthesis of organic compounds.[17] It affects the 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:[20][21]

Alpha chlorination of an aldehyde using CuCl2.

This reaction is performed in a polar solvent such as dimethylformamide, often in the presence of lithium chloride, which accelerates the reaction.[20]

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:[22]

Coupling of beta-naphthol using CuCl2.

Such compounds are intermediates in the synthesis of BINAP and its derivatives.[20]

Copper(II) chloride dihydrate promotes the hydrolysis of acetonides, i.e., for deprotection to regenerate diols[23] or aminoalcohols, as in this example (where TBDPS = tert-butyldiphenylsilyl):[24]

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 a base to give a vinyl sulfone product.[20]

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:[8]

4 HCl + O2 → 2 Cl2 + 2 H2O

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

Copper(II) chloride is used in the copper–chlorine cycle where it reacts with steam into copper(II) oxide dichloride and hydrogen chloride and is later recovered in the cycle from the electrolysis of copper(I) chloride.[11]

Niche uses

Copper(II) chloride is used in pyrotechnics as a blue/green coloring agent. In a flame test, copper chlorides, like all copper compounds, emit green-blue light.[25]

In humidity indicator cards (HICs), cobalt-free brown to azure (copper(II) chloride base) HICs can be found on the market.[26] In 1998, the European Community 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). Consequently, new cobalt-free humidity indicator cards containing copper have been developed.[27]

Copper(II) chloride is used as a mordant in the textile industry, petroleum sweetener, wood preservative, and water cleaner.[8][28]

Natural occurrence

Eriochalcite

Copper(II) chloride occurs naturally as the very rare anhydrous mineral tolbachite and the dihydrate eriochalcite.[29] Both are found near fumaroles and in some copper mines.[30][31][32] Mixed oxyhydroxide-chlorides like atacamite (Cu2(OH)3Cl) are more common, arising among Cu ore beds oxidation zones in arid climates.[33]

Safety and biological impact

Copper(II) chloride can be toxic. Only concentrations below 1.3 ppm of aqueous copper ions are allowed in drinking water by the US Environmental Protection Agency.[34] If copper chloride is absorbed, it results in headache, diarrhea, a drop in blood pressure, and fever. Ingestion of large amounts may induce copper poisoning, CNS disorders, and haemolysis.[35][36]

Copper(II) chloride has been demonstrated to cause chromosomal aberrations and mitotic cycle disturbances within A. cepa (onion) cells.[37] Such cellular disturbances lead to genotoxicity. Copper(II) chloride has also been studied as a harmful environmental pollutant. Often present in irrigation-grade water, it can negatively affect water and soil microbes.[38] Specifically, denitrifying bacteria were found to be very sensitive to the presence of copper(II) chloride. At a concentration of 0.95 mg/L, copper(II) chloride was found to cause a 50% inhibition (IC50) of the metabolic activity of denitrifying microbes.[39]

See also

References

  1. ^ A. F. Wells (1947). "The crystal structure of anhydrous cupric chloride, and the stereochemistry of the cupric atom". Journal of the Chemical Society: 1670–1675. doi:10.1039/JR9470001670.
  2. ^ Sydney Brownstein; Nam Fong Han; Eric Gabe; Yvon LePage (1989). "A redetermination of the crystal structure of cupric chloride dihydrate". Zeitschrift für Kristallographie. 189 (1): 13–15. Bibcode:1989ZK....189...13B. doi:10.1524/zkri.1989.189.1-2.13.
  3. ^ a b c NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  4. ^ Wells, A.F. (1984). Structural Inorganic Chemistry. Oxford: Clarendon Press. p. 253. ISBN 0-19-855370-6.
  5. ^ a b c d e f Greenwood, N. N. and Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. p. 1183–1185 ISBN 0-7506-3365-4.
  6. ^ Peter Baláž (2008). Mechanochemistry in Nanoscience and Minerals Engineering. Springer. p. 167. ISBN 978-3-540-74854-0.
  7. ^ Carlo Corvaja (2009). Electron paramagnetic resonance: a practitioner's toolkit. John Wiley and Sons. p. 3. ISBN 978-0-470-25882-8.
  8. ^ a b c d e f g h i j k Zhang, J.; Richardson, H. W. (2016). "Copper Compounds". Ullmann's Encyclopedia of Industrial Chemistry. pp. 1–31. doi:10.1002/14356007.a07_567.pub2. ISBN 978-3-527-30673-2.
  9. ^ Shuiliang Zhou; Shaobo Shen; Dalong Zhao; Zhitao Zhang; Shiyu Yan (2017). "Evaporation and decomposition of eutectics of cupric chloride and sodium chloride". Journal of Thermal Analysis and Calorimetry. 129 (3): 1445–1452. doi:10.1007/s10973-017-6360-y. S2CID 99924382.
  10. ^ a b c Richardson, H. W. (2003). "Copper Compounds". Kirk-Othmer Encyclopedia of Chemical Technology. doi:10.1002/0471238961.0315161618090308.a01.pub2. ISBN 0471238961.
  11. ^ a b Z. Wang; G. Marin; G. F. Naterer; K. S. Gabriel (2015). "Thermodynamics and kinetics of the thermal decomposition of cupric chloride in its hydrolysis reaction" (PDF). Journal of Thermal Analysis and Calorimetry. 119 (2): 815–823. doi:10.1007/s10973-014-3929-6. S2CID 93668361.
  12. ^ Wilhelm Biltz; Werner Fischer (1927). "Beiträge zur systematischen Verwandtschaftslehre. XLIII. Über das System Cupro-/Cuprichlorid". Zeitschrift für anorganische und allgemeine Chemie (in German). 166 (1): 290–298. doi:10.1002/zaac.19271660126.
  13. ^ A. G. Massey; N. R. Thompson; B. F. G. Johnson (1973). The Chemistry of Copper, Silver and Gold. Elsevier Science. p. 42. ISBN 9780080188607.
  14. ^ a b Naida S. Gill; F. B. Taylor (1967). Tetrahalo Complexes of Dipositive Metals in the First Transition Series. Inorganic Syntheses. Vol. 9. pp. 136–142. doi:10.1002/9780470132401.ch37. ISBN 978-0-470-13240-1.
  15. ^ a b H. Wayne Richardson (1997). Handbook of Copper Compounds and Applications. CRC Press. pp. 24–68. ISBN 9781482277463.
  16. ^ W. Libus; S. K. Hoffmann; M. Kluczkowski; H. Twardowska (1980). "Solution equilibriums of copper(II) chloride in pyridine and pyridine-diluent mixtures". Inorganic Chemistry. 19 (6): 1625–1632. doi:10.1021/ic50208a039.
  17. ^ 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. 220–223, Wiley, New York, 1738.
  18. ^ 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 978-1-85617-567-8.
  19. ^ J. Ji; W. C. Cooper (1990). "Electrochemical preparation of cuprous oxide powder: Part I. Basic electrochemistry". Journal of Applied Electrochemistry. 20 (5): 818–825. doi:10.1007/BF01094312. S2CID 95677720.
  20. ^ a b c d e f Nicholas D. P. Cosford; Pauline Pei Li; Thierry Ollevier (2015). "Copper(II) Chloride". Encyclopedia of Reagents for Organic Synthesis: 1–8. doi:10.1002/047084289X.rc214.pub3. ISBN 9780470842898.
  21. ^ C. E. Castro; E. J. Gaughan; D. C. Owsley (1965). "Cupric Halide Halogenations". Journal of Organic Chemistry. 30 (2): 587. doi:10.1021/jo01013a069.
  22. ^ 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.
  23. ^ Chandrasekhar, M.; Kusum L. Chandra; Vinod K. Singh (2003). "Total Synthesis of (+)-Boronolide, (+)-Deacetylboronolide, and (+)-Dideacetylboronolide". Journal of Organic Chemistry. 68 (10): 4039–4045. doi:10.1021/jo0269058. PMID 12737588.
  24. ^ Krishna, Palakodety Radha; G. Dayaker (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.
  25. ^ Clark, Jim (August 2018). "Flame Tests". chemguide.co.uk. Archived from the original on November 27, 2020. Retrieved January 10, 2021.
  26. ^ US 20150300958 A1, Evan Koon Lun Yuuji Hajime, "Adjustable colorimetric moisture indicators", published 2015 
  27. ^ "Cobalt dichloride". European Chemicals Agency. ECHA. Retrieved 30 May 2023.
  28. ^ B.H. Patel (2011). "11 - Natural dyes". In Clark, M. (ed.). Handbook of Textile and Industrial Dyeing. Woodhead Publishing. pp. 412–413. ISBN 9781845696955. Retrieved 2 June 2023.
  29. ^ Marlene C. Morris, Howard F. McMurdie, Eloise H. Evans, Boris Paretzkin, Harry S. Parker, and Nicolas C. Panagiotopoulos (1981) Copper chloride hydrate (eriochalcite), in Standard X-ray Diffraction Powder Patterns National Bureau of Standards, Monograph 25, Section 18; page 33.
  30. ^ "Tolbachite". mindat.org. Retrieved 24 August 2023.
  31. ^ "Eriochalcite". mindat.org. Retrieved 24 August 2023.
  32. ^ "The New IMA List of Minerals". Università degli studi di Trieste. International Mineralogical Association. Retrieved 24 August 2023.
  33. ^ "Atacamite". mindat.org. Retrieved 30 May 2023.
  34. ^ "National Primary Drinking Water Regulations". EPA. 30 November 2015. Retrieved 29 May 2023.
  35. ^ "Copper: Health Information Summary" (PDF). Environmental Fact Sheet. New Hampshire Department of Environmental Services. 2005. ARD-EHP-9. Archived from the original (PDF) on 20 January 2017.
  36. ^ "Safety Data Sheet". Sigma Aldrich. Retrieved 30 June 2023.
  37. ^ Macar, Tuğçe Kalefetoğlu (2020). "Resveratrol ameliorates the physiological, biochemical, cytogenetic, and anatomical toxicities induced by copper (II) chloride exposure in Allium cepa L." Environmental Science and Pollution Research. 27 (1): 657–667. doi:10.1007/s11356-019-06920-2. PMID 31808086. S2CID 208649491.
  38. ^ Shiyab, Safwan (2018). "Phytoaccumulation of copper from irrigation water and its effect on the internal structure of lettuce". Agriculture. 8 (2): 29. doi:10.3390/agriculture8020029.
  39. ^ Ochoa-Herrera, Valeria (2011). "Toxicity of copper (II) ions to microorganisms in biological wastewater treatment systems". Science of the Total Environment. 412 (1): 380–385. Bibcode:2011ScTEn.412..380O. doi:10.1016/j.scitotenv.2011.09.072. PMID 22030247.

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