 Dihydrate | |
_chloride.jpg) Anhydrous | |
| Names | |
|---|---|
| IUPAC name Copper(II) chloride | |
| Other names Cupric chloride | |
| Identifiers | |
| |
3D model (JSmol) |
|
| 8128168 | |
| ChEBI | |
| ChEMBL | |
| ChemSpider |
|
| DrugBank | |
| ECHA InfoCard | 100.028.373 |
| EC Number |
|
| 9300 | |
| RTECS number |
|
| UNII |
|
| UN number | 2802 |
| |
| |
| Properties | |
| CuCl2 | |
| Molar mass | 134.45 g/mol (anhydrous) 170.48 g/mol (dihydrate) |
| Appearance | dark brown solid (anhydrous) light blue 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)
|
| +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: | |
    | |
| 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) | |
| 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 | |
Otheranions | Copper(II) fluoride Copper(II) bromide |
Othercations | Copper(I) chloride Silver chloride Gold(III) chloride |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Copper(II) chloride, also known ascupric chloride, is aninorganic compound with thechemical formulaCuCl2. Themonoclinic yellowish-brownanhydrous form slowly absorbs moisture to form the orthorhombic blue-greendihydrateCuCl2·2H2O, with twowater molecules of hydration. It is industrially produced for use as aco-catalyst in theWacker process.
Both the anhydrous and the dihydrate forms occur naturally as the rare mineralstolbachite anderiochalcite, respectively.
Anhydrous copper(II) chloride adopts a distortedcadmium iodide structure. In this structure, thecopper centers areoctahedral. Most copper(II) compounds exhibit distortions from idealizedoctahedral geometry due to theJahn-Teller effect, which in this case describes the localization of oned-electron into amolecular orbital that is stronglyantibonding with respect to a pair of chloride ligands. InCuCl2·2H2O, the copper again adopts a highly distorted octahedral geometry, the Cu(II) centers being surrounded by two water ligands and four chloride ligands, whichbridge asymmetrically to other Cu centers.[4][5]
Copper(II) chloride isparamagnetic. Of historical interest,CuCl2·2H2O was used in the firstelectron paramagnetic resonance measurements byYevgeny Zavoisky in 1944.[6][7]
Aqueous solutions prepared from copper(II) chloride contain a range of copper(II)complexes depending onconcentration, temperature, and the presence of additionalchloride 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]
When copper(II) chloride solutions are treated with abase, aprecipitation ofcopper(II) hydroxide occurs:[8]
Partial hydrolysis givesdicopper 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]
Copper(II) chloride is a mildoxidant. It starts to decompose tocopper(I) chloride andchlorine gas around 400 °C (752 °F) and is completely decomposed near 1,000 °C (1,830 °F):[8][9][10][11]
The reportedmelting 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 andCuCl2.[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 withsulfur dioxide as thereductant:[8]
CuCl2 reacts with HCl or otherchloride sources to form complex ions: the red[CuCl3]− (found inpotassium trichloridocuprate(II)K[CuCl3]) (it is adimer in reality,[Cu2Cl6]2−, a couple of tetrahedrons that share an edge), and the green or yellow[CuCl4]2− (found inpotassium tetrachloridocuprate(II)K2[CuCl4]).[5][14][15]
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 ofcoordination complexes withligands such asammonia,pyridine andtriphenylphosphine oxide:[8][5][16]
However "soft" ligands such asphosphines (e.g.,triphenylphosphine), iodide, andcyanide as well as some tertiaryamines inducereduction to give copper(I) complexes.[5]
Copper(II) chloride is prepared commercially by the action ofchlorination of copper. Copper at red heat (300-400 °C) combines directly with chlorine gas, giving (molten) copper(II) chloride. The reaction is veryexothermic.[8][15]
A solution of copper(II) chloride is commercially produced by adding chlorine gas to a circulating mixture ofhydrochloric 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, orcopper(II) carbonate can react to formCuCl2 in anacid-base reaction which can subsequently be heated above 100 °C (212 °F) to produce the anhydrous derivative.[8][10]
Once prepared, a solution ofCuCl2 may be purified bycrystallization. A standard method takes the solution mixed in hot dilute hydrochloric acid, and causes the crystals to form by cooling in acalcium 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 copperelectrodes produces (among other things) a blue-greenfoam 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 generalchloralkali 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]
A major industrial application for copper(II) chloride is as a co-catalyst withpalladium(II) chloride in theWacker process. In this process,ethene (ethylene) is converted toethanal (acetaldehyde) using water and air. During the reaction,PdCl2reduced toPd, and theCuCl2 serves to re-oxidize this back toPdCl2. Air can then oxidize the resultantCuCl back toCuCl2, completing the cycle.[20]
The overall process is:[20]
Copper(II) chloride has some highly specialized applications in thesynthesis of organic compounds.[17] It affects thechlorination ofaromatic hydrocarbons—this is often performed in the presence ofaluminium oxide. It is able to chlorinate thealpha position ofcarbonyl compounds:[20][21]
This reaction is performed in a polar solvent such asdimethylformamide, often in the presence oflithium chloride, which accelerates the reaction.[20]
CuCl2, in the presence ofoxygen, can also oxidizephenols. The major product can be directed to give either aquinone or a coupled product from oxidative dimerization. The latter process provides a high-yield route to1,1-binaphthol:[22]
Such compounds are intermediates in the synthesis ofBINAP and its derivatives.[20]
Copper(II) chloride dihydrate promotes the hydrolysis ofacetonides, i.e., for deprotection to regenerate diols[23] oraminoalcohols, as in this example (where TBDPS =tert-butyldiphenylsilyl):[24]
CuCl2 also catalyses thefree radical addition ofsulfonyl chlorides toalkenes; the alpha-chlorosulfone may then undergoelimination with a base to give a vinylsulfone product.[20]
Copper(II) chloride is used as acatalyst in a variety of processes that produce chlorine byoxychlorination. TheDeacon process takes place at about 400 to 450 °C in the presence of a copper chloride:[8]
Copper(II) chloride catalyzes the chlorination in the production ofvinyl chloride anddichloromethane.[8]
Copper(II) chloride is used in thecopper–chlorine cycle where it reacts with steam into copper(II) oxide dichloride and hydrogen chloride and is later recovered in the cycle from theelectrolysis of copper(I) chloride.[11]
Copper(II) chloride is used inpyrotechnics as a blue/green coloring agent. In aflame test, copper chlorides, like all copper compounds, emit green-blue light.[25]
Inhumidity indicator cards (HICs), cobalt-free brown to azure (copper(II) chloride base) HICs can be found on the market.[26] In 1998, theEuropean Community classified items containing cobalt(II) chloride of 0.01 to 1%w/w as T (Toxic), with the correspondingR 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 amordant in the textile industry,petroleumsweetener,wood preservative, andwater cleaner.[8][28]
Copper(II) chloride occurs naturally as the very rare anhydrous mineral tolbachite and the dihydrate eriochalcite.[29] Both are found nearfumaroles and in some copper mines.[30][31][32] Mixed oxyhydroxide-chlorides likeatacamite (Cu2(OH)3Cl) are more common, arising among Cu ore beds oxidation zones in arid climates.[33]
Copper(II) chloride can be toxic. Only concentrations below 1.3ppm of aqueous copper ions are allowed in drinking water by theUS Environmental Protection Agency.[34] If copper chloride is absorbed, it results in headache, diarrhea, a drop inblood pressure, and fever. Ingestion of large amounts may inducecopper poisoning,CNS disorders, andhaemolysis.[35][36]
Copper(II) chloride has been demonstrated to causechromosomal aberrations andmitotic cycle disturbances withinA. cepa (onion) cells.[37] Such cellular disturbances lead togenotoxicity. 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]
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