Copper(II) chloride

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Copper(II) chloride
Tolbachite-3D-balls.png
Anhydrous
Copper(II) chloride.jpg
Anhydrous
Cupric chloride.jpg
Dihydrate
Names
Other names
Cupric chloride
Identifiers
3D model (JSmol)
8128168
ChEBI
ChEMBL
ChemSpider
DrugBank
ECHA InfoCard 100.028.373 OOjs UI icon edit-ltr-progressive.svg
EC Number
  • 231-210-2
9300
PubChem CID
RTECS number
  • GL7000000
UNII
UN number 2802
  • InChI=1S/2ClH.Cu/h2*1H;/q;;+2/p-2 Yes check.svgY
    Key: ORTQZVOHEJQUHG-UHFFFAOYSA-L Yes check.svgY
  • 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)
Appearanceyellow-brown solid (anhydrous)
blue-green solid (dihydrate)
Odor odorless
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)
Solubility methanol:
68 g/100 mL (15 °C)


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

+1080·10−6 cm3/mol
Structure
distorted CdI2 structure
Octahedral
Hazards
GHS labelling:
GHS-pictogram-acid.svg GHS-pictogram-skull.svg GHS-pictogram-exclam.svg GHS-pictogram-pollu.svg
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)
2
0
1
Flash point Non-flammable
NIOSH (US health exposure limits):
PEL (Permissible)
TWA 1 mg/m3 (as Cu) [1]
REL (Recommended)
TWA 1 mg/m3 (as Cu) [1]
IDLH (Immediate danger)
TWA 100 mg/m3 (as Cu) [1]
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).
X mark.svgN  verify  (what is  Yes check.svgYX mark.svgN ?)
Copper(II) chloride dihydrate Dikhlorid medi digidrat.jpg
Copper(II) chloride dihydrate
Copper(II) chloride anhydrous Dikhlorid medi bezvodnyi.jpg
Copper(II) chloride anhydrous

Copper(II) chloride is the chemical compound with the chemical formula CuCl2. The anhydrous form is yellowish brown but slowly absorbs moisture to form a blue-green dihydrate.

Contents

Both the anhydrous and the dihydrate forms occur naturally as the very rare minerals tolbachite and eriochalcite, respectively. [2]

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. [3]

Copper(II) chloride is paramagnetic. Of historical interest, CuCl2·2H2O was used in the first electron paramagnetic resonance measurements by Yevgeny Zavoisky in 1944. [4] [5]

Properties and reactions

Aqueous solutions of copper(II) chloride. Greenish when high in [Cl ], more blue when lower in [Cl ]. CuCl2 equilibrium.JPG
Aqueous solutions of copper(II) chloride. Greenish when high in [Cl ], more blue when lower in [Cl ].

Depending on the concentration, temperature, and presence of extra chloride ions, copper(II) chloride-prepared aqueous solutions can contain a variety of copper(II) complexes. 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−. [6]

Hydrolysis

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

CuCl2 + 2 NaOH → Cu(OH)2 + 2 NaCl
Copper(II) chloride dihydrate crystal Copper(II)chloride crystal 01.jpg
Copper(II) chloride dihydrate crystal

Partial hydrolysis gives dicopper chloride trihydroxide, Cu2(OH)3Cl, a popular fungicide.

Redox

Copper(II) chloride is a mild oxidant. It decomposes to copper(I) chloride and chlorine gas near 1000 °C:

2 CuCl2 → 2 CuCl + Cl2

Copper(II) chloride (CuCl2) 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:

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 (it is a dimer in reality, Cu2Cl62−, a couple of tetrahedrons that share an edge), and the green or yellow CuCl42−. [7]

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

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

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

CuCl2 + 2 C5H5N → [CuCl2(C5H5N)2] (tetragonal)
CuCl2 + 2 (C6H5)3PO → [CuCl2((C6H5)3PO)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.

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.

Cu(s) + Cl2(g) → CuCl2(l)

It is also commercially practical to combine copper(II) oxide with an excess of ammonium chloride at similar temperatures, producing copper chloride, ammonia, and water:[ citation needed ]

CuO + 2NH4Cl → CuCl2 + 2NH3 + H2O

Although copper metal itself cannot be oxidised 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.

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. [8] [9]

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.

Natural occurrence

Copper(II) chloride occurs naturally as the very rare anhydrous mineral tolbachite and the dihydrate eriochalcite. [2] Both are found near fumaroles and in some Cu mines. [10] [11] [12] 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

In organic synthesis

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

Other organic synthetic applications

Copper(II) chloride has some highly specialized applications in the synthesis of organic compounds. [8] It affects 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: [13]

CuCl2 alpha chlorination.png

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: [14]

CuCl2 naphthol coupling.png

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 [15] or aminoalcohols, as in this example (where TBDPS = tert-butyldiphenylsilyl): [16]

CuCl2 DeprotectionOfAminoAlcohol.png

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 ]

In inorganic synthesis

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 catalyzes the chlorination in the production of vinyl chloride and dichloroethane. [17]

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.

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.

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.

Safety

Copper(II) chloride can be toxic. Only concentrations below five ppm are allowed in drinking water by the US Environmental Protection Agency.[ citation needed ]

Copper (II) chloride has been demonstrated to cause chromosomal aberrations and mitotic cycle disturbances within A. cepa cells. [18] Such cellular disturbances lead to genotoxicity. Copper (II) chloride has also been studied as a harmful environmental pollutant. Often present in irrigation-grade water, copper (ii) chloride can negatively affect water and soil microbes. [19] Specifically, denitrifying bacteria were found to be very sensitive to the presence of copper(ii) chloride. At a concentration of 0.95 mg L−1, copper(II) chloride was found to cause a 50% inhibition (IC50) of the metabolic activity of denitrifying microbes. [20]

Related Research Articles

<span class="mw-page-title-main">Chlorine</span> Chemical element, symbol Cl and atomic number 17

Chlorine is a chemical element with the symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the revised Pauling scale, behind only oxygen and fluorine.

<span class="mw-page-title-main">Sodium hypochlorite</span> Chemical compound (known in solution as bleach)

Sodium hypochlorite is an inorganic chemical compound with the formula NaOCl, comprising a sodium cation and a hypochlorite anion. It may also be viewed as the sodium salt of hypochlorous acid. The anhydrous compound is unstable and may decompose explosively. It can be crystallized as a pentahydrate NaOCl·5H
2
O
, a pale greenish-yellow solid which is not explosive and is stable if kept refrigerated.

Iron(III) chloride is the inorganic compound with the formula FeCl3. Also called ferric chloride, it is a common compound of iron in the +3 oxidation state. The anhydrous compound is a crystalline solid with a melting point of 307.6 °C. The colour depends on the viewing angle: by reflected light the crystals appear dark green, but by transmitted light they appear purple-red.

<span class="mw-page-title-main">Copper(II) nitrate</span> Chemical compound

Copper(II) nitrate describes any member of the family of inorganic compounds with the formula Cu(NO3)2(H2O)x. The hydrates are blue solids. Anhydrous copper nitrate forms blue-green crystals and sublimes in a vacuum at 150-200 °C. Common hydrates are the hemipentahydrate and trihydrate.

<span class="mw-page-title-main">Lead(II) chloride</span> Chemical compound

Lead(II) chloride (PbCl2) is an inorganic compound which is a white solid under ambient conditions. It is poorly soluble in water. Lead(II) chloride is one of the most important lead-based reagents. It also occurs naturally in the form of the mineral cotunnite.

Neodymium(III) chloride or neodymium trichloride is a chemical compound of neodymium and chlorine with the formula NdCl3. This anhydrous compound is a mauve-colored solid that rapidly absorbs water on exposure to air to form a purple-colored hexahydrate, NdCl3·6H2O. Neodymium(III) chloride is produced from minerals monazite and bastnäsite using a complex multistage extraction process. The chloride has several important applications as an intermediate chemical for production of neodymium metal and neodymium-based lasers and optical fibers. Other applications include a catalyst in organic synthesis and in decomposition of waste water contamination, corrosion protection of aluminium and its alloys, and fluorescent labeling of organic molecules (DNA).

<span class="mw-page-title-main">Aluminium chloride</span> Chemical compound

Aluminium chloride, also known as aluminium trichloride, is an inorganic compound with the formula AlCl3. It forms hexahydrate with the formula [Al(H2O)6]Cl3, containing six water molecules of hydration. Both are colourless crystals, but samples are often contaminated with iron(III) chloride, giving a yellow color.

<span class="mw-page-title-main">Manganese(II) chloride</span> Chemical compound

Manganese(II) chloride is the dichloride salt of manganese, MnCl2. This inorganic chemical exists in the anhydrous form, as well as the dihydrate (MnCl2·2H2O) and tetrahydrate (MnCl2·4H2O), with the tetrahydrate being the most common form. Like many Mn(II) species, these salts are pink, with the paleness of the color being characteristic of transition metal complexes with high spin d5 configurations.

<span class="mw-page-title-main">Cobalt(II) chloride</span> Chemical compound

Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl
2
. The compound forms several hydrates CoCl
2
·nH
2
O
, for n = 1, 2, 6, and 9. Claims of the formation of tri- and tetrahydrates have not been confirmed. The anhydrous form is a blue crystalline solid; the dihydrate is purple and the hexahydrate is pink. Commercial samples are usually the hexahydrate, which is one of the most commonly used cobalt compounds in the lab.

<span class="mw-page-title-main">Copper(I) chloride</span> Chemical compound

Copper(I) chloride, commonly called cuprous chloride, is the lower chloride of copper, with the formula CuCl. The substance is a white solid sparingly soluble in water, but very soluble in concentrated hydrochloric acid. Impure samples appear green due to the presence of copper(II) chloride (CuCl2).

<span class="mw-page-title-main">Chromium(III) chloride</span> Chemical compound

Chromium(III) chloride (also called chromic chloride) describes any of several chemical compounds with the formula CrCl3 · xH2O, where x can be 0, 5, and 6. The anhydrous compound with the formula CrCl3 is a violet solid. The most common form of the trichloride is the dark green hexahydrate, CrCl3 · 6 H2O. Chromium chlorides find use as catalysts and as precursors to dyes for wool.

<span class="mw-page-title-main">Nickel(II) chloride</span> Chemical compound

Nickel(II) chloride (or just nickel chloride) is the chemical compound NiCl2. The anhydrous salt is yellow, but the more familiar hydrate NiCl2·6H2O is green. Nickel(II) chloride, in various forms, is the most important source of nickel for chemical synthesis. The nickel chlorides are deliquescent, absorbing moisture from the air to form a solution. Nickel salts have been shown to be carcinogenic to the lungs and nasal passages in cases of long-term inhalation exposure.

<span class="mw-page-title-main">Iron(II) chloride</span> Chemical compound

Iron(II) chloride, also known as ferrous chloride, is the chemical compound of formula FeCl2. It is a paramagnetic solid with a high melting point. The compound is white, but typical samples are often off-white. FeCl2 crystallizes from water as the greenish tetrahydrate, which is the form that is most commonly encountered in commerce and the laboratory. There is also a dihydrate. The compound is highly soluble in water, giving pale green solutions.

<span class="mw-page-title-main">Wacker process</span>

The Wacker process or the Hoechst-Wacker process refers to the oxidation of ethylene to acetaldehyde in the presence of palladium(II) chloride as the catalyst. This chemical reaction was one of the first homogeneous catalysis with organopalladium chemistry applied on an industrial scale.

<span class="mw-page-title-main">Gold(III) chloride</span> Chemical compound

Gold(III) chloride, traditionally called auric chloride, is a compound of gold and chlorine with the molecular formula Au2Cl6. The "III" in the name indicates that the gold has an oxidation state of +3, typical for many gold compounds. Gold(III) chloride is hygroscopic and decomposes in visible light. This compound is a dimer of AuCl3. This compound has few uses, although it catalyzes various organic reactions.

<span class="mw-page-title-main">Tin(II) chloride</span> Chemical compound

Tin(II) chloride, also known as stannous chloride, is a white crystalline solid with the formula SnCl2. It forms a stable dihydrate, but aqueous solutions tend to undergo hydrolysis, particularly if hot. SnCl2 is widely used as a reducing agent (in acid solution), and in electrolytic baths for tin-plating. Tin(II) chloride should not be confused with the other chloride of tin; tin(IV) chloride or stannic chloride (SnCl4).

Calcium hypochlorite is an inorganic compound with formula Ca(OCl)2. It is a white solid, although commercial samples appear yellow. It strongly smells of chlorine, owing to its slow decomposition in moist air. This compound is relatively stable as a solid and solution and has greater available chlorine than sodium hypochlorite. "Pure" samples have 99.2% active chlorine. Given common industrial purity, an active chlorine content of 65-70% is typical. It is the main active ingredient of commercial products called bleaching powder, used for water treatment and as a bleaching agent.

<span class="mw-page-title-main">Metal halides</span>

Metal halides are compounds between metals and halogens. Some, such as sodium chloride are ionic, while others are covalently bonded. A few metal halides are discrete molecules, such as uranium hexafluoride, but most adopt polymeric structures, such as palladium chloride.

Cobalt(III) chloride or cobaltic chloride is an unstable and elusive compound of cobalt and chlorine with formula CoCl
3
. In this compound, the cobalt atoms have a formal charge of +3.

<span class="mw-page-title-main">Copper compounds</span> Chemical compounds containing copper

Copper forms a rich variety of compounds, usually with oxidation states +1 and +2, which are often called cuprous and cupric, respectively. Copper compounds, whether organic complexes or organometallics, promote or catalyse numerous chemical and biological processes.

References

  1. 1 2 3 NIOSH Pocket Guide to Chemical Hazards. "#0150". National Institute for Occupational Safety and Health (NIOSH).
  2. 1 2 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.
  3. Wells, A.F. (1984) Structural Inorganic Chemistry, Oxford: Clarendon Press. ISBN   0-19-855370-6.
  4. Peter Baláž (2008). Mechanochemistry in Nanoscience and Minerals Engineering. Springer. p. 167. ISBN   978-3-540-74854-0.
  5. Marina Brustolon (2009). Electron paramagnetic resonance: a practitioner's toolkit. John Wiley and Sons. p. 3. ISBN   978-0-470-25882-8.
  6. Greenwood, N. N. and Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN   0-7506-3365-4.
  7. 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.
  8. 1 2 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, 1738.
  9. 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.
  10. "Tolbachite".
  11. "Eriochalcite".
  12. "List of Minerals". 21 March 2011.
  13. C. E. Castro; E. J. Gaughan; D. C. Owsley (1965). "Cupric Halide Halogenations". Journal of Organic Chemistry . 30 (2): 587. doi:10.1021/jo01013a069.
  14. 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.
  15. 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.
  16. Krishna, Palakodety Radha; G. Dayaker (2007). "A stereoselective total synthesis of (−)-andrachcinidine via an olefin cross-metathesis protocol". Tetrahedron Letters . Elsevier. 48 (41): 7279–7282. doi:10.1016/j.tetlet.2007.08.053.
  17. H.Wayne Richardson, "Copper Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, 2005, Wiley-VCH, Weinheim, doi : 10.1002/14356007.a07_567
  18. 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.
  19. 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 .
  20. 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.

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.