Copper(I) chloride

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Copper(I) chloride
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Template:Longitem 8127933
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Template:Longitem CuCl
Molar mass 98.999 g/mol[1]
Appearance white powder, slightly green from oxidized impurities
Density 4.14 g/cm3[1]
Melting point Template:Chembox CalcTemperatures
Boiling point Template:Chembox CalcTemperatures
Template:Longitem 1.72×10−7
Solubility insoluble in ethanol,
acetone;[1] soluble in concentrated HCl, NH4OH
Band gap 3.25 eV (300 K, direct)[2]
Template:Longitem −40.0·10−6 cm3/mol[3]
Template:Longitem 1.930[4]
Template:Longitem Zincblende, cF20
Template:Longitem F43m, No. 216[5]
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a = 0.54202 nm
NFPA 704 (fire diamond) Template:NFPA 704 diamond
Flash point Template:Chembox CalcTemperatures
Template:Longitem Copper(I) fluoride
Copper(I) bromide
Copper(I) iodide
Template:Longitem Silver(I) chloride
Gold(I) chloride
Template:Longitem Copper(II) chloride

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File:IR Spectrum of Copper(I) chloride.png
IR absorption spectrum of copper(I) chloride

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).

History

Copper(I) chloride was first prepared by Robert Boyle and designated rosin of copper in the mid-seventeenth century from mercury(II) chloride ("Venetian sublimate") and copper metal:[7]

HgCl2 + 2 Cu → 2 CuCl + Hg

In 1799, Joseph Proust first differentiated two different chlorides of copper. He prepared CuCl (which he called white muriate of copper) by heating CuCl2 at red heat in the absence of air, causing it to lose half of its combined chlorine followed by removing residual CuCl2 by washing with water.[8]

An acidic solution of CuCl was formerly used to analyze carbon monoxide content in gases, for example in Hempel's gas apparatus where the CuCl absorbs the carbon monoxide.[9] This application was significant during the nineteenth and early twentieth centuries when coal gas was widely used for heating and lighting.[10]

Synthesis

Copper(I) chloride is produced industrially by the direct combination of copper metal and chlorine at 450–900 °C:[11][12]

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Copper(I) chloride can also be prepared by reducing copper(II) chloride with sulfur dioxide, or with ascorbic acid (vitamin C) that acts as a reducing sugar:[13][14]

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Many other reducing agents can be used.[12]

Properties

Copper(I) chloride has the cubic zincblende crystal structure at ambient conditions. Upon heating to 408 °C the structure changes to hexagonal. Several other crystalline forms of CuCl appear at high pressures (several GPa).[5]

Upon contact with water, copper(I) chloride slowly undergoes disproportionation:[15]

2 CuCl → Cu + CuCl2

In part for this reason, samples in air assume a green coloration.[16]

Copper(I) chloride is a Lewis acid. It is classified as soft according to the hard-soft acid-base concept. Thus, it forms a series of complexes with soft Lewis bases such as triphenylphosphine:

CuCl + 1 P(C6H5)3 → 1/4 {CuCl[P(C6H5)3]}4
CuCl + 2 P(C6H5)3 → CuCl[P(C6H5)3)]2
CuCl + 3 P(C6H5)3 → CuCl[P(C6H5)3)]3

CuCl also forms complexes with halides. For example H3O+ CuCl2− forms in concentrated hydrochloric acid.[17] Chloride is displaced by CN− and S2O32−.[12]

Solutions of CuCl in HCl absorb carbon monoxide to form colourless complexes such as the chloride-bridged dimer [CuCl(CO)]2. The same hydrochloric acid solutions also react with acetylene gas to form [CuCl(C2H2)]. Ammoniacal solutions of CuCl react with acetylenes to form the explosive copper(I) acetylide, Cu2C2. Alkene complexes of CuCl can be prepared by reduction of CuCl2 by sulfur dioxide in the presence of the alkene in alcohol solution. Complexes with dienes such as 1,5-cyclooctadiene are particularly stable:[18]

Structure of COD complex of CuCl

Uses

The main use of copper(I) chloride is as a precursor to the fungicide copper oxychloride. For this purpose aqueous copper(I) chloride is generated by comproportionation and then air-oxidized:[12]

Cu + CuCl2 → 2 CuCl
4 CuCl + O2 + 2 H2O → Cu3Cl2(OH)4 + CuCl2

Copper(I) chloride catalyzes a variety of organic reactions, as discussed above. Its affinity for carbon monoxide in the presence of aluminium chloride is exploited in the COPureSM process.[19]

In organic synthesis

CuCl is used as a co-catalyst with carbon monoxide, aluminium chloride, and hydrogen chloride in the Gatterman-Koch reaction to form benzaldehydes.[20]

In the Sandmeyer reaction, the treatment of an arenediazonium salt with CuCl leads to an aryl chloride. For example:[21][22]

(Example Sandmeyer reaction using CuCl)

The reaction has wide scope and usually gives good yields.[22]

Early investigators observed that copper(I) halides catalyse 1,4-addition of Grignard reagents to alpha,beta-unsaturated ketones[23] led to the development of organocuprate reagents that are widely used today in organic synthesis:[24]

(Addition of RMgX to C=C-C=O mediated by CuCl)

This finding led to the development of organocopper chemistry. For example, CuCl reacts with methyllithium (CH3Li) to form "Gilman reagents" such as (CH3)2CuLi, which find use in organic synthesis. Grignard reagents form similar organocopper compounds. Although other copper(I) compounds such as copper(I) iodide are now more often used for these types of reactions, copper(I) chloride is still recommended in some cases:[25]

(Alkylation of sorbate ester at 4-position mediated by CuCl)

Cuprous chloride also catalyzes the dimerization of acetylene to vinylacetylene, once used as a precursor to various polymers such a neoprene.[26]

Niche uses

CuCl is used as a catalyst in atom transfer radical polymerization (ATRP). It is also used in pyrotechnics as a blue/green coloring agent. In a flame test, copper chlorides, like all copper compounds, emit green-blue.[27]

Natural occurrence

Natural form of CuCl is the rare mineral nantokite.[28][29]

References

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  1. ↑ a b c d e Template:RubberBible92nd
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  4. ↑ Patnaik, Pradyot (2002) Handbook of Inorganic Chemicals. McGraw-Hill, Template:ISBN
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  16. ↑ Pastor, Antonio C. (1986) U.S. patent 4582579 "Method of preparing cupric ion free cuprous chloride" Section 2, lines 4–41.
  17. ↑ Script error: No such module "Citation/CS1".
  18. ↑ Nicholls, D. (1973) Complexes and First-Row Transition Elements, Macmillan Press, London.
  19. ↑ Script error: No such module "Citation/CS1".
  20. ↑ Script error: No such module "Citation/CS1".
  21. ↑ Wade, L. G. (2003) Organic Chemistry, 5th ed., Prentice Hall, Upper Saddle River, New Jersey, p. 871. Template:ISBN.
  22. ↑ a b March, J. (1992) Advanced Organic Chemistry, 4th ed., Wiley, New York. p. 723. Template:ISBN
  23. ↑ Script error: No such module "Citation/CS1".
  24. ↑ Jasrzebski, J. T. B. H.; van Koten, G. (2002) Modern Organocopper Chemistry, N. Krause (ed.). Wiley-VCH, Weinheim, Germany. p. 1. Script error: No such module "CS1 identifiers". Template:ISBN.
  25. ↑ Bertz, S. H.; Fairchild, E. H. (1999) Handbook of Reagents for Organic Synthesis, Volume 1: Reagents, Auxiliaries and Catalysts for C-C Bond Formation, R. M. Coates, S. E. Denmark (eds.). Wiley, New York. pp. 220–3. Template:ISBN.
  26. ↑ Script error: No such module "Citation/CS1".
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External links

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